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.
4. Two Major Types of
Bonding
• Ionic Bonding
– forms ionic compounds
– transfer of e-
• Covalent Bonding
– forms molecules
– sharing e-
5. 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)
6. It’s the mobile electrons
that enable me-tals to
conduct electricity!!!!!!
7. 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
8. IONic bonding
• Always formed between metals and
non-metals
[METALS ]+ [NON-METALS ]-
Lost e-
Gained e-
9. 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.
11. • 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
12. Covalent Bonding
• Pairs of e- are shared
between non-metal atoms
• electronegativity difference < 2.0
• forms polyatomic ions
molecules
13. 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
14. Covalent, Ionic, metallic
bonding?
• NO2
• sodium
hydride
• Hg
• H2S
• sulfate
• NH4
+
• Aluminum
phosphate
• KH
• KCl
• HF
• CO
• Co
Also study
your
characteristics!
15. Drawing ionic compounds
using Lewis Dot Structures
• Symbol represents the KERNEL of the
atom (nucleus and inner e-)
• dots represent valence e-
16. NaCl
• This is the finished Lewis Dot
Structure
[Na]+ [ Cl ]
-
How did we get here?
17. • 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
19. Drawing molecules using
Lewis Dot Structures
• Symbol represents the KERNEL of the
atom (nucleus and inner e-)
• dots represent valence e-
20. 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
21. Methane CH4
• This is the finished Lewis dot structure
How did we get here?
22. • 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!
23. 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
24. • DOUBLE bond
– atoms that share two e- pairs (4 e-)
O O
• TRIPLE bond
– atoms that share three e- pairs (6 e-)
N N
25. Draw Lewis Dot Structures
You may represent valence electrons
from different atoms with the
following symbols x, ,
CO2
NH3
26. 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!!!!!
28. Types of Covalent Bonds
• NON-Polar bonds
–Electrons shared evenly in the bond
–E-neg difference is zero
Between identical atoms
Diatomic molecules
29. 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”
30. 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
31. Polar molecules (a.k.a.
Dipoles)
• Not equal on all sides
–Polar bond between 2 atoms makes a
polar molecule
–asymmetrical shape of molecule
35. Making sense of the polar
non-polar thing
BONDS
Non-polar Polar
Identical Different
MOLECULES
Non-polar Polar
Symmetrical Asymmetrical
36. IONIC bonds ….
Ionic bonds are
so polar that the electrons are not
shared but transferred between
atoms forming ions!!!!!!
37. C. Johannesson
VSEPR Theory
• Valence Shell Electron Pair Repulsion
Theory
• Electron pairs orient themselves in
order to minimize repulsive forces.
38. C. Johannesson
VSEPR Theory
• Types of e- Pairs
– Bonding pairs - form bonds
– Lone pairs - nonbonding e-
Lone pairs repel
more strongly than
bonding pairs!!!
44. • Attractions between
molecules
– van der Waals forces
• Weak attractive
forces between
non-polar
molecules
– Hydrogen “bonding”
• Strong attraction
between special
polar molecules
Intermolecular attractions
45. 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
46. van der Waals periodicity
• increase with molecular mass.
• increase with closer distance between
molecules
– Decreases when particles are farther away
47. Hydrogen “Bonding”
• Strong polar
attraction
– Like magnets
• Occurs ONLY
between H of one
molecule and N, O,
F of another
H “bond”
48. H is shared between
2 atoms of OXYGEN or
2 atoms of NITROGEN or
2 atoms of FLUORINE
Of
2
different
molecules
49. Why does H “bonding”
occur?
• Nitrogen, Oxygen and Fluorine
– small atoms with strong nuclear charges
• powerful atoms
– very high electronegativities
50. Intermolecular forces
dictate chemical properties
• Strong intermolecular forces cause
high b.p., m.p. and slow evaporation
(low vapor pressure) of a substance.
51. 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
