The document outlines the characteristics and formation of ionic bonds, emphasizing the electrostatic attraction between oppositely charged ions and factors such as electronegativity and periodic table positions. It also discusses covalent bonds, molecular geometry through VSEPR theory, and intermolecular forces like hydrogen bonds. Additionally, it addresses properties like solubility, melting points, and electrical conductivity of compounds.

1. Bonding
Topic 4

2. Describe the ionic bond

3. The ionic bond
• Electrostatic attraction between
oppositely charged ions.
• Elements in groups 1-7 (Analogy).
• Transition elements can form more than
one ion (Fe+2, Fe+3).
• Some compounds exist as ions
(Polyatomic ions).

4. Ionic bond formation
• Ionization process.
Energy
Na
Cl
Metal atom
Non-metal atom
Na+ Cl-
Electrostatic
attraction
Sodium chloride

5. In order to form an ionic compound
1. Position in the Periodic Table
!
• DOWN A GROUP, metals tend to lose
electrons.
• UP A GROUP, non-metals tend to gain
electrons.
• So, ionic bonds usually are formed
between atoms far from each other.

7. 2. Electronegativity
!
• If the difference in electronegativity equals to 1.8
or higher than 1.8, the bond is ionic.
!
• Data Booklet.
Na2O
MgO
Al2O3
SiO2
P2O5

8. Lattice structures
• 3D structures.
!
• Lattice enthalpy.
!
• Geometry depends on size of ions.
!
• Coordination number (NaCl is 6).

9. • Ionic structures http://
www.chemguide.co.uk/atoms/structures/
ionicstruct.html#top
!
• Jmol: http://www.chemeddl.org/resources/
models360/solids.php#nacl

10. Practice
1. Solve some problems in pages 112 and
114.
!
2. Class will be divided in 4 groups; each
one will build and describe an ionic
molecule. Once they are built, the class
will discuss them and order them in
increasing ionic character.
!
3. HW #2: Solve problems on Edmodo.

11. Covalent bonds

12. http://www.mhhe.com/physsci/chemistry/
animations/chang_7e_esp/bom1s2_11.swf

13. Lewis structure
• Represents the valence shell.
!
• Ethane, hydroxide ion, ammonium, ethyne,
CF3Cl.
!
• Incomplete or expanded octate.
BeCl2 and BF3.
They tend to form dative bonds (CO).

14. Strength of a bond
• What affects the strength?
!
• Electrons can spend more time in the
neighbouring atom.
!
Polar bonds (assymetrical bond).
Higher than 0 and smaller than 1.8
Non-polar = 0

15. VSEPR theory
• Valence Shell Electron Pair Repulsion.
!
• Deduce geometry.
!
• Importance.
!
• Negative charge center.

16. VSEPR theory
• http://www.mhhe.com/physsci/chemistry/
animations/chang_7e_esp/bom3s2_7.swf

17. Shapes
!
!
http://phet.colorado.edu/en/simulation/
molecule-shapes

18. Polarity based on shape
• Net pull = 0. Non-polar
• Net pull greater than 0. Polar
!
CCl4
CH3Cl
H2O

19. Polarity based on shape
• http://www.mhhe.com/physsci/chemistry/
animations/chang_7e_esp/bom4s2_7.swf

20. Allotropes
http://www.chemeddl.org/resources/
models360/solids.php#diamond

21. Diamond Graphite Fullerene
•Repeated
sequence of
tetrahedrals.
•Hardest
natural
substance.
•sp3 C-C.
•Does not
conduct
•Hexagon
parallel layers.
•sp2 C-C.
•van der Waals´
forces between
layers (weak).
•Good
conductor.
•12 pentagons
and 20
hexagons.
•sp2 C-C.
•Sphere.
•Semiconductor.

22. Si and SiO2
Si
• Tetrahedral.
• Giant (covalent) lattice.
!
SiO2
• Quartz.
• Giant covalent.
• Each Si attached to 4 O, and each O attached
to 2 Si.
!
http://www.chemeddl.org/resources/models360/
solids.php#quartz

23. Shapes of molecules and ions

24. Five negative charges
• Triangular bipyramidal.

25. Five negative charges
• One one-bonding pair.
• Seesaw.
• 180, 117.

26. Five negative charges
• Two non-bonding pairs.
• T-shaped.
• 90,180.

27. Five negative charges
• Three non-bonding pairs.
• 180.
• Linear.
• Triiodide ion

28. Six negative charges
• Octahedral.
• 90.

29. Six negative charges
• One non-bonding pair.
• Square pyramidal.
• 180, 90.

30. Six negative charges
• Two non-bonding pair.
• Square planar.
• 180.

31. Hybridization
Sigma bonds (simple) and Pi bonds (double-triple)
!
• http://www.mhhe.com/physsci/chemistry/
essentialchemistry/flash/hybrv18.swf
!
Hybridizations and geometry
!
• http://www.mhhe.com/physsci/chemistry/
animations/chang_7e_esp/bom5s2_6.swf

34. Intermolecular forces
• van der Waals’ forces
Non-polar molecules.
Temporary dipole.
Induced dipole.
Increases as size increases.

35. Intermolecular forces
• Dipole-dipole attraction
Polar molecules have permanent dipole.
When two molecules meet (same) a dipole-
dipole attraction occurs.

36. Intermolecular forces
• Hydrogen bonds
Hydrogens HAS to be attached to O, F or
N.

38. Quick question
Which compound will have the highest
boiling point, Water or Hydrofluoric
acid?
!
WATER

39. Metallic bonding
• Electrons are delocalized (current).
• The remaining cations form a lattice.
“Lattice metal cations in a sea of delocalized
electrons”.

40. • Melleability: shaped.
!
!
!
• Ductility: Threads.

41. Physical properties
1. Solubility
• How easy can a solute disperse a
solvent.
• Polar dissolves polar.
• H-bonds are also responsible.
!
2. Melting-boiling point
• Ionic compounds tend to have high
melting points because of the lattice.

42. 3. Electrical conductivity
!
• Depends on movile ions.
• Covalent compounds do not conduct
electricity, except some, like HCl
(Electronegativity difference = 1.2)

43. THANKYOU FORYOUR
ATTENTION