Download Chemical Bonds ch6 p.161

This is one of the
chapters you must
read….
chapter 6…bonding
Student will learn:
1. three types of bonding
ionic, covalent, metallic
2. two categories of bonding
polar, non-polar
3. how to draw Lewis structure
4. how to calculate electronegative
5. bond characteristics
6. VSPER theory
Chemical Bonds
ch6 p.161
?What holds chemicals together?
Chemical Bonds:
electrical attraction between +nuclei and
–valence electron of different atoms
• By bonding together the atoms are more
stable, and have a lower level of energy
arrangement
3 types of Bonding:
l. Ionic Bonding: lose or gain –e
metal + Non-metal = ionic bonding
= makes ions
share –e
Non-metal + non-metal=covalent bonds
= makes molecules
2. Covalent Bonding:
3. Metallic Bonding: -e flow free in a sea of –e
Transition metals
Pick the bonding
NaCl, CH4, HCl, K2S, FeSO4, LiF, H20, Cu, Zn, Mg(OH)2
2 catagories for the bonding
polar
non-polar
Unequal
attraction for electrons
equal
balanced attraction
Ionic bonding (Metals+nonMetals) is always polar
Covalent (nonMetals +nonMetals) maybe either polar/nonpolar
2 ways to figure out
Draw Lewis structure
calculate and use chart
iodomethane
CH3I
1. Lewis Dot for each element
C
H
2. Arrange to form skeleton
If a carbon then always in middle
Lewis Structure for
I
If no carbon then least electronegative atom in
middle
Hydrogen never in middle
See that it is lopsided……
polar covalent molecule
How about individual bonds?
Lewis for ammonia
NH3
1. Draw Lewis dot for each element
N
H
2. Do skeleton : hydrogen never in middle
Does it look lopsided….
polar covalent molecule
Lewis for formaldehyde
l. Lewis dot for each
CH20
C
H
O
2. Do skeleton: carbon always in middle
Notice left out –e….move to make a double bond
Single bonds
Double bonds
Triple bonds
Notice lopsided: polar……..covalent
Lewis struture for :
CCl4
l. Draw Lewis dot for each:
C
2. Draw skeleton: carbon always in middle
Does it look lopsided?........
No…..
non-polar covalent molecule
What about each bond?
CL
Use table on p.151 and chart on page 162
Calculating Polarity of Bonds:
using
differences in electronegativity
Remember: electronegativity = ability to gain electrons
Bonding is rarely purely ionic or covalent………most of time somewhere in
between
Use table on p.151 and chart on page 162 (overhead 31)
Subtract the two electronegativity numbers then ? is it less than 1.7
= polar covalent?
Calculate bond type and polarity
KCl,
MgCl2,
H2,
H2S
Cs2S,
SCL2,
Comparing Characteristics
Ionic bonds (vs) Covalent bonds
Metals + nonmetals
Gain or lose electrons so….
+ or – ends ……very polar
Will form a crystalline lattice,
look on page 177 (model)
Stronger bonds
Most are solids
Higher melting point
Higher boiling point
Many dissolve in water,
+ion and -ion break apart in water so will conduct
electricity in water. Some do not dissolve because
the pull between the charges are greater than
the attraction of H2O molecule
Hard but brittle----why?
A shift of one row of ions causes a large build
up of repulsive forces. And --do not like-so if one layer moves that forces the other
layers to move so they are brittle.
Non-metals + Non-metals
Share electrons
Exist as individual molecules
Weaker bonds
Most are gases, some liquids
Very low melting point
Very low boiling point
Will evaporate at room
temperature
Overhead 70
Ionic compounds
form
Crystalline lattice
How ionic compounds dissolve
Why Ionic Compounds are brittle
Comparing Characteristics
Ionic bonds (vs) Covalent bonds
Metals + nonmetals
Gain or lose electrons so….
+ or – ends ……very polar
Will form a crystalline lattice,
look on page 177 (model)
Stronger bonds
Most are solids
Higher melting point
Higher boiling point
Many dissolve in water,
+ion and -ion break apart in water so will conduct
electricity in water. Some do not dissolve because
the pull between the charges are greater than
the attraction of H2O molecule
Hard but brittle----why?
A shift of one row of ions causes a large build
up of repulsive forces. And --do not like-so if one layer moves that forces the other
layers to move so they are brittle.
Non-metals + Non-metals
Share electrons
Exist as individual molecules
Weaker bonds
Most are gases, some liquids
Very low melting point
Very low boiling point
Will evaporate at room
temperature
Overhead 70
Why most
covalents
are liquids
or gases
and
evaporate
easy.
Metallic Bonding p.181
Transitional Metals: vacant outer p orbitals
because filling up d orbitals first.
4s2, 3d10, 4p…
they overlap
This overlapping lets –e roam freely about the metal
network of empty atomic orbitals.
These mobile –e form a sea of electrons which are
packed in a lattice form.
Overhead 68
Characteristics of Metallic bonding “Cu,Au,Ag, Fe”
l.
Conduct electricity
Conduct heat ::::::
due to the “sea of electrons”
ability to move freely
2.
Reflect light, Shiny, Polish
:::::: Contain many orbitals (d10) separated
by extremely small energy differences, metals can absorb a
wide range of light frequencies. This absorption of light
energy accounts for the ability to reflect light and be shiny.
…..p. 181
3.
Malleable: hammer into a thin sheet.
::::::possible because the metallic bonding
is same in all directions throughout the solid
because of “sea of electrons”
Ductile: ability to be drawn into a thin wire.
::::::Because the metallic bonding is same in
all directions throughout the solid because of
“sea of electrons”
VSEPR THEORY
Valence Shell Electron Pair Repulsion
Theory : replusion between Valence Shell Electrons Pairs
surrounding an atom causes these sets to be oriented as
“far apart as possible”.
“AS FAR APART AS POSSIBLE”
Lewis dot, VSEPR TO PREDICT GEOMETRY OF MOLECULE,
Intermolecular force:
3 types:
the attraction between molecules
dipole-dipole
hydrogen bonding
London dispersion forces
Dipole-dipole: strongest intermolecular force
created when by equal but opposite charges of the
molecule come within close distance of each other.
Hydrogen bonding:
hydrogen atom that is bonded to a
highly electronegative atom is attracted to an unshared pair of
electrons of an electronegative atom in a near-by molecule.
“this is why water expands when frozen”
London Dispersion force: results
from the constant motion of electrons and
the creation of instantaneous dipoles.
Because London dispersion force depends
on the motion of electrons, the strength
increases with increasing atomic masses.