Chapter
6:
Chemical
Bonding
Cartoon courtesy of NearingZero.net
6-1 Chemical Bonds
• Forces that hold groups of atoms
together and make them function as a
unit.
• Ionic bonds – transfer of electrons
Metal + Nonmetal
Ex) NaCl Li2O
• Covalent bonds – sharing of electrons.
2 nonmetals
Ex) H2O CO2
• Metallic bonds- electrons are free to
move throughout the material. Metals
Covalent Bonds
1. Nonpolar-Covalent bonds (H2)
• Electrons are equally shared
• Electronegativity difference of 0 to
0.3
2. Polar-Covalent bonds (HCl)
• Electrons are unequally shared
• Electronegativity difference between .3
and 1.7
Using Electronegativity
differences
• SCs
• S=2.5
• Cs= 0.7
2.5-0.7= 1.8
look on table pg 162
1.7 thru 3.3= Ionic Bond
6-2 Covalent Bonding
• Molecule- is the smallest unit quanitity of
matter which can exist by itself and retain
all the properties of the original
substance.
• Examples H2O & O2
• Diatomic molecule- is a molecule containing
2 identical atoms. (H2 N2 O2 F2 Cl2 Br2 I2)
• H NO F
• Chemical Formula- represents the
relative # of atoms of each kind in a
chemical compound by using atomic
symbols and numerical subscripts.
Example: H2O H=2 O=1
• Molecular compound (Covalent
Compounds) - simplest formula unit are
molecules. Have low melting & boiling pts.
• Molecular formula- shows the types and
numbers of atoms combined in a single
molecule.
• Bond Length- is the average
distance between 2 bonded atoms.
• Bond Energy- is the energy required
to break a bond.
• It gives us information about the
strength of a bonding interaction.
• Resonance-Occurs when more than
one valid Lewis structure can be
written for a particular molecule.
The Octet Rule
• Chemical
compounds tend to
form so that each
atom, by gaining,
losing, or sharing
electrons, has an
(8) octet of
electrons in its
valence shell.
H
He
8 is
Great!
Lewis Dot
• Shows how valence
electrons are
arranged among
atoms in a molecule.
• Reflects central idea
that stability of a
compound relates to
noble gas electron
configuration.
I
started
it for
you
CH4
Try one yourself NH3
H2O
Structural Formula
• Shows shared pair
of electrons by a
dashed line.
•
•
•
•
•
•
Single bond- 1 pair of electrons
Double bond- 2 pair of electrons
Triple bond- 3 pair of electrons
Try a couple:
O2
N2
6-3 Ionic Bonding and
Compounds
• Ionic Compound- is composed of
positive and negative ions combined
so that the positive & negative
charges are equal. (Metal + nonmetal)
• Formula Unit- is the simplest
collection of atoms from which a
compounds formula can be
established.
1.
2.
3.
Writing Formulas
Write the symbols of each element
Put their charge in their upper right corner
Crisscross the numbers down (Not the charges).
Example:
Write the formula for Magnesium Chloride
Mg Cl
Mg+2 Cl-1
Mg+2 Cl-1
MgCl
MgCl22
Writing Formulas Practice
+1
•
1.
2.
3.
4.
+2
Write the formula for:
Aluminum Bromide
Calcium Oxide
Calcium Nitride
Sodium Chloride
+3
4 -3 -2 -1
• Lattice Energy- is the energy
released when 1 mole of an ionic
crystalline compound is formed from
gaseous ions.
• Ionic Compounds have high melting
points & boiling points, are hard and
brittle, have crystalline structure.
Polyatomic Ions
• Many atoms with a
charge.
• Example SO4-2
6-4 Metallic Bonding
• Metals- conduct heat, have low
ionization energy & electronegativity,
give up e• Metallic Bond- is a chemical bond
resulting from the attraction
between positive ions and surrounding
mobile e-.
• Malleability and ductility
6-5 Molecular Geometry
• VSEPR Theory- “Valence- shell,
electron-pair repulsion”
• states that repulsion between the
sets of valence-level electrons
surrounding an atom cause these sets
to be oriented as far apart as
possible.
Determining VSEPR
•
1.
2.
3.
4.
5.
Determine the VSEPR
for H2O
Draw the Lewis Dot
Draw the Structural
Formula
Label the central atom
as A
Label any atoms
attached to the center
atom as B
Label any paired
electrons on the central
atom that are not used
in the bond as E
H-O-H
B
E2
A B
VSEPR AB2E2
Shape Bent (look on chart)
VSEPR Chart
VSEPR
AB or AB2
AB2E
AB3
AB4
AB3E
AB2E2
AB5
AB6
SHAPE
Linear
Bent
Trigonal-Planar
Tetrahedral
Trigonal-Pyramidal
Bent
Trigonal-Bipyramidal
Octahedral
• Hybridization-The Blending of Orbitals.
• Dipole- is created by equal but opposite charges
that are separated by a short distance.
• Dipole-Dipole Attractions-Attraction between
oppositely charged regions of neighboring
molecules.
• Hydrogen Bonding- Bonding between hydrogen
and more electronegative neighboring atoms such
as oxygen and nitrogen. Hydrogen bonding in
Kevlar, a strong polymer used in bullet-proof
vests.
• London Dispersion Forces- The temporary
separations of charge that lead to the London
force attractions are what attract one nonpolar
molecule to its neighbors. London forces increase
with the size of the molecules.
•
•
•
•
•
•
•
“Electronegativity chart”. Table. Aug. 9, 2006.
http://www.chemistry210.com/notes/u01s06f.htm
“Lewis Structures”. Drawings. Aug. 9, 2006. http://www.avonchemistry.com/chem_bond_explain.html
“Oscar”. Photo. Aug. 9, 2006. http://www.musicmerchant.com/22061.htm
“Water Structural Formula”. Drawing. Aug. 10, 2006.
http://www.accs.net/users/kriel/ch4notes/water_structura
l_formula.gif
“Periodic Table of Elements”. Chart. Aug. 9, 2006.
http://users.erols.com/kdennis/periodictable.jpg
“Information”. Aug 11, 2006.
http://www.sciencegeek.net/Chemistry/Powerpoin
t/Unit3/Unit3_files/frame.htm
Holt, Rinehart and Winston. Modern Chemistry. Harcourt Brace & Company.
1999.