ib chem chemistry bonding science atoms stick ions ionic covalent metal metallic #awesome

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