1. HC
CHEMISTRY HC
CHEMISTRY
(B) Periodicity
Bonding Continuum

2. HC HC Bonding Continuum (B) Periodicity
CHEMISTRY HC
CHEMISTRY
(B) Periodicity
Bonding Continuum
After completing this topic you should be able to :
Learners should be familiar with ionic and covalent bonding. In a covalent bond, atoms share pairs of electrons.
The covalent bond is a result of two positive nuclei being held together by their common attraction for the shared pair of electrons.
Polar covalent bonds are formed when the attraction of the atoms for the pair of bonding electrons is different. Delta positive (δ + ) and delta negative (δ- ) notation can be used to indicate the partial charges on atoms, which give rise to a dipole (eg H δ+ ̶ Clδ- ).
Pure covalent bonding and ionic bonding can be considered as being at opposite ends of a bonding continuum with polar covalent bonding lying between these two extremes. If the difference is large then the movement of bonding electrons from the element of lower electronegativity to the element of higher electronegativity is complete resulting in the formation of ions. Compounds formed between metals and non-metals are often, but not always ionic.

3. Polar covalent bonds Learning intention
Learn how differences in electronegativity between bonding atoms lead to the formation of polar covalent bonds.

4. Polar Covalent Bonds
Non-polar covalent bond – electrons shared equally between atoms (same electronegativity)
Polar covalent bond – electrons shared unequally between atoms (atom B is more electronegative)

5. Covalent Bonding
A covalent bond is a shared pair of electrons electrostatically attracted to the positive nuclei of two atoms.
Both nuclei try to pull the electrons towards themselves - + - +
This is like a tug-of-war where both sides are pulling on the same object.
It creates a strong bond between the two atoms.
The atoms achieve a stable outer electron arrangement (a noble gas arrangement) by sharing electrons.

6. Covalent Bonding Picture a tug-of-war:
If both teams pull with the same force the mid-point of the rope will not move.

7. Pure Covalent Bond H e H
This even sharing of the rope can be compared to a pure covalent bond, where the bonding pair of electrons are held at the mid-point between the nuclei of the bonding atoms.

8. Covalent Bonding What if it was an uneven tug-of-war?
The team on the right are far stronger, so will pull the rope harder and the mid-point of the rope will move to the right.

9. Polar Covalent Bond
A polar covalent bond is a bond formed when the shared pair of electrons in a covalent bond are not shared equally.
This is due to different elements having different electronegativities.

10. Polar Covalent Bond I δ- δ+ H e.g. Hydrogen Iodide e
If hydrogen iodide contained a pure covalent bond, the electrons would be shared equally as shown above.
This makes iodine slightly negative and hydrogen slightly positive. This is known as a dipole.
However, iodine has a higher electronegativity and pulls the bonding electrons towards itself
(winning the tug-of-war)

11. Polar Covalent Bond C Cl δ+ δ-
In general, the electrons in a covalent bond are not equally shared. δ- δ+
e.g. C Cl
2.5
3.0
Electronegativities
δ- indicates where the bonding electrons are most likely to be found.

12. Polar Covalent Bond Consider the polarities of the following bonds:
Electronegativities
Difference
C Cl
0.5
P H
O H
1.3
C Cl δ- δ+
O H δ- δ+
P H
Increasing Polarity
Complete a similar table for C-N, C-O and P-F bonds.

13. Polar Covalent Bonds
In the covalent bond between fluorine and hydrogen. The bonding electrons are not shared equally between the two atoms.
Hydrogen
Fluorine
The fluorine nucleus has more protons and has a stronger pull on the electrons than the hydrogen nucleus..

14. -
Thus the fluorine atom has a greater share of the bonding electrons and acquires a slight negative charge. + F H
The hydrogen atom is then made slightly positive.
The bond is a polar covalent bond and we use the symbols + and - to show this.
The dipole produced is permanent.
Fluorine is the most electronegative element. It is small atom compared to others and its nucleus is massive for its atomic size.
Some other polar covalent bonds are O-H and N-H

15. Permanent dipole-permanent dipole interactions
Learning intention
Learn about this additional intermolecular force of attraction which exists between polar molecules.

16. Dipole-Dipole Attractions
The differing electronegativities of different atoms in a
molecule and the spatial arrangement of polar covalent bonds
can cause a molecule to form a permanent dipole. H
Permanent dipole
Asymmetrical molecule + Cl Cl - Cl - -
3 polar covalent C–Cl bonds and
1 polar covalent C-H bond in CHCl3
POLAR molecule

17. Permanent Dipole-Dipole interactions
The attraction is stronger than Londons forces
Hydrogen bonding is a particular example of dipole-dipole attractions.
Molecules with permanent dipoles attract each other.

18. Bond Strengths Bond Type Strength (kJ mol –1) Metallic 80 to 600 Ionic
Covalent
Hydrogen 40
Dipole-Dipole 30
London’s forces
1 to 20

19. Polar molecules and permanent dipoles
Both propanone and butane have the same formula mass of 58
however, butane boils at – 1 oC while propanone boils at 56oC
Propanone is a polar molecule as it has a permanent dipole, so has polar-polar attraction as well as London’s forces between molecules. - +
b.p. 56 o C
Butane has no permanent dipoles, so only London’s forces
between molecules. So has a lower boiling point.
b.p. -1 o C

20. 4 polar covalent C-Cl bonds in CCl4 tetrahedral shape
Symmetry CCl4
No permanent dipole
Symmetrical molecule Cl - + Cl Cl - - Cl
Tetrachloromethane has a symmetrical arrangement of polar bonds and the polarity cancels out over the molecule. -
4 polar covalent C-Cl bonds in CCl4 tetrahedral shape
NON-POLAR molecule

21. 2 polar covalent C=O bonds in CO2 linear shape
Symmetry CO2
No permanent dipole
Symmetrical molecule + - - O O
2 polar covalent C=O bonds in CO2 linear shape
Carbon dioxide has a symmetrical arrangement of polar bonds and the polarity cancels out over the molecule.
NON-POLAR molecule

22. Bonding continuum Learning intention
Learn about the bonding continuum which stretches between pure covalent and ionic bonds, in terms of differences of electronegativity between bonding atoms.

23. Electronegativity Difference and Bond Type:
Difference Bond Example
Covalent (nonpolar) H-H 0.0
Covalent (polar) H-Cl H20 0.7
Covalent (very polar) H-F 1.9
 2.0 Ionic NaCl 2.1

24. The greater the difference in electronegativity the greater the polarity between two bonding atoms and the more ionic in character.

25. Bonding Continuum “Covalent compounds are formed by non-metals only”
IS NOT AN ABSOLUTE LAW!
Some compounds break this rule….

26. Making Tin(IV)iodide
Gently heat the tin and iodine in a small conical flask containing 10cm3 of tolulene on a hot plate.
Collect the yellow precipitate by filtration using Büchner filtration

27. Making Tin(IV)iodide
Determine the melting point of the solid collected.

28. Making Tin(IV)iodide Melting point of tin(IV)iodide is 143oC.
Tin electronegativity of 1.8
Iodine has electronegativity of 2.6
Molecule contains polar covalent bonds, but the symmetry cancels out the dipoles, therefore only weak London dispersion forces so low melting an boiling point.

29. Titanium (IV) chloride
TiCl4 is a dense, colourless liquid.
It is one of the rare transition metal halides that is a liquid at room temperature, This property reflects the fact that TiCl4 is ………….; that is, each TiCl4 ………. is relatively …………… associated with its neighbours.
Most metal chlorides are ionic. The attraction between the individual TiCl4 molecules is weak, primarily ……………….……….., and ……………. these weak Van der Waals (intermolecular) interactions result in low melting and boiling points.
TiCl4 is soluble in toluene and dichloromethane, as are other non-polar species.

30. TiCl4
Used in smoke grenades and for smoke screens