This document discusses different types of chemical bonds including ionic bonds, covalent bonds, and metallic bonds. Ionic bonds form when a metal transfers electrons to a non-metal. Covalent bonds form when two non-metals share electrons. Metallic bonds form through electrostatic attraction between positively charged metal ions and delocalized electrons. Examples of different bonds are also provided such as lithium fluoride forming an ionic bond through Li+ and F- ions and metallic bonding in metals occurring through interaction of positively charged atomic residues and free-flowing electrons.

1. Chemical bonding
and Lewis structure

2. A force that holds
atoms together to
create a single unit.
Chemical bond

3. Metallic
Bonds
Metal +
Metal
Covalent
Bonds
Non-Metal +
Non-Metal
Electrons are
shared between
atoms.
Ionic
Bonds
Metal + Non-
Metal
Electrons are
transferred.

4. Valence electrons
• Electrons in the outer energy level
of atom.
• Are involved in chemical bonding.
• Are easily removed.
Noble gases obey the octet structure.

5. Lewis Diagrams
Electron pair Unpaired electron

6. When two non-metallic
atoms bond together by
sharing one or more pairs
of electrons.
Covalent bond

7. •Forms molecules.
•Molecule is held together by a
shared electron pair.
•Individual atoms in the
molecule has octet structure.

8. Covalent bond – single bond.

9. Covalent bond – single bond.

10. Covalent bond – double bond

11. Covalent bond – triple bond

12. Relative molecular mass is
the sum of the relative
atomic mass of the atoms
in the molecule.

13. Results from the transfer
of one or more electrons
from metal atoms to non-
metal atoms.
Ionic bond

14. Metals give off electrons to form cations in order to
obtain a noble gas electron structure.
Li  Li+ + e-
[He]2s1 [He]
[Ar]4s2 [Ar]
Ca Ca2++2e-

15. Non-metals gain electrons to form anions in order to
obtain a noble gas electron structure.
N + 3e-  [ N ]3-
[He]2s2p3
S + 2e-  [ S ]2-
[Ne]3s23p4
[He]2s22p6 = [Ne]
[Ne]3s23p6 = [Ar]

16. • The metal atom gives away its electrons to the
non-metal.
• The positive and negative ions attract each
other due to electrostatic forces. (opposite
charges attract)
• The electrostatic attraction is the ionic bond.
• Crystal lattices form, and are built up of
alternating positive and negative ions.

17. Ca  Ca2+ + 2e-
2x: F + e-  [ F ]-
Ca2+ + 2[ F ]-  Ca2+2[ F ]-  CaF2
Positive ion:
Negative ion:
Crystal lattice

18. Now you do:
Lithium Fluoride Li+ and F-
Aluminium sulfide Al3+ and S2-

19. Ionic compounds are in fact composed of
many millions of ions, and not just one
cation and one anion. This huge structure is
known as a crystal lattice.
NaCl is the FORMULA unit (not a molecule
of salt.)

21. Ionic bonds form between
metals and radicals.
Radicals will also form ionic
bonds with NH4
+

22. The sum of the relative
atomic masses of the
atom in the chemical
formula.
Relative formula mass

23. Electrostatic attraction
between the positive
atomic residue and the
sea of delocalised
electrons. Metallic bonds

25. • Atoms have 1, 2, or 3 valence electrons.
• Metallic atoms have a low ionisation energy.
• Little energy required to remove valence electrons –
electrons delocalise.
• Atoms are tightly packed to form a metal lattice
where the valence orbital overlap.
• Atoms that lose an electron becomes positive (called
a positive atomic residue).
• The bonding energy is the electrostatic force of
attraction between the sea of delocalised electrons,
and the positive atomic residue.

26. Properties of
metals
Reason
Metal glow Sea of delocalised electrons can reflect light and
therefore causes the surface to shine.
Electrical conductor Sea of delocalised electrons can move freely and act as
charge carriers.
Thermal conductor Sea of delocalised electrons act as carriers of heat energy
(kinetic energy).
Malleable/Ductile Atoms, though firmly bound together, can slide over
each other, allowing metal to bend and stretch.
High density Atoms in the solid phase is packed closely together in a
metal lattice.

27. Alloys

28. Metals are often converted into
alloys to change their properties.
Alloys are usually always stronger
than pure metals.

29. Text: Olivier, A: Physical Science textbook and workbook, Grade 10 ; Reivilo Uitgewers.
Slide 8 – 11: Lily Kotzè
Slide 18: Lily Kotzè
Slide 20: http://dtc-wsuv.org/isci/module/img/7_14_NaCl.png
Slide 24: http://i.imgur.com/oIoEZhH.gif
Slide 27: http://i0.wp.com/upload.wikimedia.org/wikipedia/commons/thumb/3/36/Alloy_Substitutional.svg/150px-
Alloy_Substitutional.svg.png?w=625
Slide 28: Steel: https://encrypted-
tbn2.gstatic.com/images?q=tbn:ANd9GcRb7xMBVsBuBLTvAOfjtLXoghcMRgMkIH5puPfnzKRRKmwUNtHk
Stainless steel: http://curiousscience.com/images/stock/img780.jpg Bronze:
https://upload.wikimedia.org/wikipedia/commons/c/ca/Apa_Schwerter.jpg