1. Chemical Bonding
Why do bonds form?
to lower the potential energy
between positive and
negative charges
positive charges
protons
cations
negative charges
electrons
anions
non-metals
metals
gain e-
lose e-
Periodic Table

2. Lewis electron-dot symbols
element symbol
= nucleus + core e-
one “dot” =
valence e-
metals
dot = e- it loses to form cation
non-metal
unpaired dot =
e- paired through
e- gain or sharing

3. Ionic bonding
metal
+ non-metal
Na (s) 2
+ Cl2 (g) 
NaCl (s) : .
+ Cl- : Na
+ Cl : .
Na+ :
[Ne]
3s1
[Ne]
3s23p5
[Ne]
[Ar] 2
Ca(s)
+ O2(g)  2
CaO (s) . . . Ca : O : .
[Ar]4s2
[He]2s2 2p4
Ca2+
O2-
[Ar]
[Ne]

4. high Electron Affinity
Ionic bonding
metal
+ non-metal
low Ionization Energy
high Electron Affinity
lose 1 or 2 valence e-
gain e-
electron transfer takes place
electrostatic attraction between cation and anion e- - +
formula
= ratio of anions to cations

5. Ionic sizes isoelectronic series same # electrons 46 e-+
isoelectronic series
same # electrons
46 e-
+49
+50
+51
ions get smaller

6. Ionic bonding
metal
+ non-metal 2
Na (s)
+ Cl2 (g)  2
NaCl (s)

8. Ionic bonding
metal
+ non-metal 2
Na (s)
+ Cl2 (g)  2
NaCl (s)
exothermic
heat given off
negative
Ionization Energy Na
Na+
+ 496 kJ/mol
Electron Affinity Cl
Cl-
-349 kJ/mol
Lattice Energy
E = k Q1 Q2
-787 kJ/mol d
-640 kJ/mol
Coulomb’s law
Na+
Cl-
NaCl +

9. Ionic solids cation + anion + - metal non-metal sodium + chlorine NaCl
chloride Na Cl + -
801o C
lithium +
oxygen
Li O
lithium
oxide Li O 2 + 2-
> 1700oC
magnesium +
nitrogen
Mg N
magnesium
nitride Mg N 2+ 3- 3 2
strong interactions
(ion-ion)
high melting points

10. Transition metals
more than 1 form
except
Ag+
Zn2+
Cd2+
Al3+
aluminum
manganese +
oxygen +
sulfur Al
Mn1+
Mn2+
Mn3+
Mn4+ S 3+ 2-
Al S 2 3
aluminum sulfide
Mn3+ O 2-
Mn2O3
Mn4+
O2-
MnO2
manganese(IV) oxide
manganese(III) oxide

11. Covalent bonding
non-metal
+ non-metal
electrons shared between atoms
high Ionization Energies
high Electron Affinities
electron density
between the atoms
distance between atoms =
bond length
formula =
actual # atoms

12. Lewis structure : . : F : . : F . F : F + :
[He]2s22p5
[Ne]
e- not used in bonding
lone pairs
shared e-
bonding pair
shared equally between F : . . . H O : : O . H
1s1
[He]2s22p4
[He]
[Ne]
oxygen
2 lone pairs
bonding pair
not shared equally

13. Covalent bonding
non-metal
+ non-metal H . H . H . +
1s1
1s1
[He]

14. share valence electrons = chemical bonds carbon + chlorine C Cl 4+ -
Covalent compounds
non-metal +
non-metal
share valence electrons
= chemical bonds
carbon +
chlorine
mono di
tri
tetra
penta
hexa
hepta
octa C Cl 4+ -
CCl4
carbon
tetra
chloride
nitrogen +
oxygen N ? O 2- NO
nitrogen monoxide 2+
NO2
nitrogen dioxide 4+
N2O4
dinitrogen
tetroxide 4+

15. Covalent compounds N H
NH3
nitrogen
trihydride
ammonia 3+ 1- H O
H2O
dihydrogen
monoxide
water 1+ 2-
weak forces
low m.p.
0.0oC
Table 2.5 p. 62
Polyatomic ions
NH4+
ammonium
ClO3-
chlorate
OH-
hydroxide
MnO4-
permanganate
NO3-
nitrate
CrO42-
chromate
SO42-
sulfate
CO32-
carbonate
PO43-
phosphate

16. Metallic Bonding
metal
+ metal
metals
valence e-
well shielded
low Ionization Energy
low Electron Affinities
share valence e-
not localized between atoms
delocalized
move freely throughout metal Na
(nucleus and core e-)
e- “sea”
(valence e-)

17. Electronegativity
ability of an atom
in a molecule
to attract e- to itself
related to
Ionization Energy
Electron Affinity
Pauling scale

18. NaCl 2.1 ionic 801oC BeCl2 1.5 polar covalent 405oC AlCl3 1.5
PCl3
0.9
polar covalent
76oC
Cl2
0.0
covalent
non-polar covalent  0.4 C H
polar covalent + C O -
ionic
> 1.8
Li2O