Chemical Bonding

1. Chemical
BONDING

2. Chemical Bond
• A bond results from the attraction of nuclei
for electrons
– All atoms trying to achieve a stable octet
• IN OTHER WORDS
– the p+ in one nucleus are attracted to the e- of
another atom
• Electronegativity

3. Two Major Types of
Bonding
• Ionic Bonding
– forms ionic compounds
– transfer of e-
• Covalent Bonding
– forms molecules
– sharing e-

4. One minor type of bonding
• Metallic bonding
– Occurs between like atoms of a metal in the
free state
– Valence e- are mobile (move freely among all
metal atoms)
– Positive ions in a sea of electrons
• Metallic characteristics
– High mp temps, ductile, malleable, shiny
– Hard substances
– Good conductors of heat and electricity as (s) and (l)

5. It’s the mobile electrons
that enable me-tals to
conduct electricity!!!!!!

6. IONic Bonding
• electrons are transferred between
valence shells of atoms
• ionic compounds are
made of ions
• ionic compounds are called Salts or
Crystals
NOT MOLECULES

7. IONic bonding
• Always formed between metals and
non-metals
[METALS ]+ [NON-METALS ]-
Lost e-
Gained e-

8. IONic Bonding
• Electronegativity difference > 2.0
– Look up e-neg of the atoms in the bond
and subtract
NaCl
CaCl2
• Compounds with polyatomic ions
NaNO3

10. • hard solid @ 22oC
• high mp temperatures
• nonconductors of electricity in solid
phase
• good conductors in liquid phase or
dissolved in water (aq)
SALTS
Crystals
Properties of Ionic
Compounds

11. Covalent Bonding
• Pairs of e- are shared
between non-metal atoms
• electronegativity difference < 2.0
• forms polyatomic ions
molecules

12. Properties of Molecular
Substances
• Low m.p. temp and b.p. temps
• relatively soft solids as compared
to ionic compounds
• nonconductors of electricity in
any phase
Covalent
bonding

13. Covalent, Ionic, metallic
bonding?
• NO2
• sodium
hydride
• Hg
• H2S
• sulfate
• NH4
+
• Aluminum
phosphate
• KH
• KCl
• HF
• CO
• Co
Also study
your
characteristics!

14. Drawing ionic compounds
using Lewis Dot Structures
• Symbol represents the KERNEL of the
atom (nucleus and inner e-)
• dots represent valence e-

15. NaCl
• This is the finished Lewis Dot
Structure
[Na]+ [ Cl ]
-
How did we get here?

16. • Step 1 after checking that it is IONIC
– Determine which atom will be the +ion
– Determine which atom will be the - ion
• Step 2
– Write the symbol for the + ion first.
• NO DOTS
– Draw the e- dot diagram for the – ion
• COMPLETE outer shell
• Step 3
– Enclose both in brackets and show each charge

17. Draw the Lewis Diagrams
• LiF
• MgO
• CaCl2
• K2S

18. Drawing molecules using
Lewis Dot Structures
• Symbol represents the KERNEL of the
atom (nucleus and inner e-)
• dots represent valence e-

19. Always remember atoms are
trying to complete their
outer shell!
The number of electrons the atoms
needs is the total number of bonds
they can make.
Ex. … H? O? F? N? Cl? C?
one two one three one four

20. Methane CH4
• This is the finished Lewis dot structure
How did we get here?

21. • Step 1
– count total valence e- involved
• Step 2
– connect the central atom (usually the first in
the formula) to the others with single bonds
• Step 3
– complete valence shells of outer atoms
• Step 4
– add any extra e- to central atom
IF the central atom has 8 valence e- surrounding
it . . YOU’RE DONE!

22. Sometimes . . .
• You only have two atoms, so there is
no central atom, but follow the same
rules.
• Check & Share to make sure all the
atoms are “happy”.
Cl2 Br2 H2 O2 N2 HCl

23. • DOUBLE bond
– atoms that share two e- pairs (4 e-)
O O
• TRIPLE bond
– atoms that share three e- pairs (6 e-)
N N

24. Draw Lewis Dot Structures
You may represent valence electrons
from different atoms with the
following symbols x, ,
CO2
NH3

25. Draw the Lewis Dot Diagram for
polyatomic ions
• Count all valence e- needed for
covalent bonding
• Add or subtract other electrons based
on the charge
REMEMBER!
A positive charge means it LOST
electrons!!!!!

26. Draw Polyatomics
• Ammonium
• Sulfate

27. Types of Covalent Bonds
• NON-Polar bonds
–Electrons shared evenly in the bond
–E-neg difference is zero
Between identical atoms
Diatomic molecules

28. Types of Covalent Bonds
Polar bond
–Electrons unevenly shared
–E-neg difference greater than zero
but
less than 2.0
closer to 2.0 more polar
more “ionic character”

29. non-polar MOLECULES
• Sometimes the bonds within a
molecule are polar and yet the
molecule is non-polar because its
shape is symmetrical. H
H
H
H C
Draw Lewis dot first and
see if equal on all sides

30. Polar molecules (a.k.a.
Dipoles)
• Not equal on all sides
–Polar bond between 2 atoms makes a
polar molecule
–asymmetrical shape of molecule

31. H Cl -
+

32. H
H
O

-
+
Water is asymmetrical
+

33. Water is a bent molecule
O
H H H H

34. Making sense of the polar
non-polar thing
BONDS
Non-polar Polar
Identical Different
MOLECULES
Non-polar Polar
Symmetrical Asymmetrical

35. IONIC bonds ….
Ionic bonds are
so polar that the electrons are not
shared but transferred between
atoms forming ions!!!!!!

36. C. Johannesson
VSEPR Theory
• Valence Shell Electron Pair Repulsion
Theory
• Electron pairs orient themselves in
order to minimize repulsive forces.

37. C. Johannesson
VSEPR Theory
• Types of e- Pairs
– Bonding pairs - form bonds
– Lone pairs - nonbonding e-
Lone pairs repel
more strongly than
bonding pairs!!!

38. 4 Shapes of molecules

39. 1. Linear (straight line)
Ball and stick
model
Space filling
model

40. 2. Bent
Ball and stick
model
Space filling
model

41. 3.Trigonal pyramid
Ball and stick
model
Space filling
model

42. 4.Tetrahedral
Ball and stick
model
Space filling
model

43. • Attractions between
molecules
– van der Waals forces
• Weak attractive
forces between
non-polar
molecules
– Hydrogen “bonding”
• Strong attraction
between special
polar molecules
Intermolecular attractions

44. van der Waals
• Non-polar molecules can exist in liquid
and solid phases
because van der Waals forces keep the
molecules attracted to each other
• Exist between CO2, CH4, CCl4, CF4,
diatomics and monoatomics

45. van der Waals periodicity
• increase with molecular mass.
• increase with closer distance between
molecules
– Decreases when particles are farther away

46. Hydrogen “Bonding”
• Strong polar
attraction
– Like magnets
• Occurs ONLY
between H of one
molecule and N, O,
F of another
H “bond”

47. H is shared between
2 atoms of OXYGEN or
2 atoms of NITROGEN or
2 atoms of FLUORINE
Of
2
different
molecules

48. Why does H “bonding”
occur?
• Nitrogen, Oxygen and Fluorine
– small atoms with strong nuclear charges
• powerful atoms
– very high electronegativities

49. Intermolecular forces
dictate chemical properties
• Strong intermolecular forces cause
high b.p., m.p. and slow evaporation
(low vapor pressure) of a substance.

50. Which substance has the
highest boiling point?
• HF
• NH3
• H2O
• WHY?
Fluorine has the highest e-neg,
SO
HF will experience the
strongest H bonding and 
needs the most energy to
weaken the i.m.f. and boil

51. The End