Transcript Chapter 8

Chapter 8
Chemical Bonding
Octet Rule
Octet Rule: Atoms will gain, lose, or share
valence electrons in order to obtain a stable
outer shell electron configuration of 8 valence
electrons. (H and He only need 2)
Lewis Electron Dot Structures
36
4
7
X
2
1
58
Lewis Structures – use dots to
represent valence electrons
Lewis Electron Dot Structures
Another way
Draw the Lewis Structure for:
1. Phosphorus
2. Chlorine
3. Magnesium Ion
4. Oxygen Ion
Chemical Bonds
1. Chemical Bonds: Attractive forces that hold
atoms together due to the mutual attraction
between the nuclei and their valence
electrons.
2. Three types of Chemical Bonds
Ionic
Covalent (Polar & Nonpolar)
Metallic
Metallic Bonds
Metallic bonds: result from the delocalized
sharing of valence electrons between metal
atoms.
Example: Aluminum, Copper, Iron
Electronegativity
Electronegativity – the ability of an atom in a
molecule to attract electrons to itself
Metals = low values
(Ionization energy too)
Nonmetals = high values
(Ionization energy too)
Note Trend:
Across  increase
Down  decrease
Types of Bonds
Ionic
Type 
Electronegativity
Difference…
Valence
Electrons are…
How to
Recognize…
Chemical Formula
for one example…
Polar Nonpolar
Covalent Covalent
>1.7
1.7>x>0
= zero
(equal to or more
than 1.7)
(less than 1.7 but
more than zero)
Transferred
Shared
unequally
Shared
Equally
Metal &
Nonmetal
2 Different
Nonmetals
2 of the same
Nonmetals
NaCl
H2O
H2
Ionic Bonding
1. Metals lose electrons
and form cations . Na +
2. Nonmetals gain
electrons and form
anions. Cl –
3. Ionic Bonds form due to
the attractions between
the metal ions and
nonmetal ions.
Lattice Energy
1. Lattice Energy is the energy needed to separate
an ionic lattice into gaseous ions.
2. Lattice Energy increases with:
 Increasing charge on the ions
LE stronger in MgO than NaCl
 Decreasing distance between the ions
(LE is stronger between smaller radii ions)
LE stronger in LiF than in CsBr
3. The greater the LE, the higher the melting point.
Explain These Trends in Lattice Energy
1. NaCl > RbBr
2. BaO > KF
3. CaF2 > BaF2
HW - Textbook
#8.1
#8.8
#8.12ab
#8.14
#8.19
#8.21
#8.22
Covalent Bonds vs Ionic Bonds
Covalent Bonds: form when valence electrons are
shared between atoms of 2 nonmetals
Ionic Bonds: form when valence electrons are
transferred from a metal atom to a nonmetal atom
Covalent Bonds
1. Single Bond – sharing of 1 electron pair
2. Double Bond – sharing of 2 electron pairs
3. Triple Bond – sharing of 3 electron pairs
Bond Length & Strength
Single bonds: Longer Bond Lengths
Lowest Bond Energies (weak bonds)
Triple bonds: Shorter Bond Lengths
Highest Bond Energies (strong bonds)
Bond Polarity
1. Nonpolar Covalent Bond – electrons are
shared equally between two atoms
Br2
2. Polar Covalent Bond – one atom exerts a
greater attraction for the bonding atoms
H2O
3. The greater the electronegativity difference
between the atoms, the more polar the
bond.
Which Bond is Most Polar?
C–O
or Si – F
B – Br or
S - Cl
Rules for Drawing Lewis Structures for
Molecules & Polyatomic Ions
1. Find the number of valence electrons in each
atom and add them up.
(Note: - ions have gained e-, + ions have lost)
2. Draw symbols of the atoms near each other in
the way they will bond.
3. Connect the atoms by single bonds (1 e- pair)
4. Complete octets on atoms bonded to the central
atom. (except hydrogen gets 2)
5. Place leftover electrons on the central atom.
6. If there aren’t enough electrons to give the
central atom an octet, try double or triple
bonds.
Do These
Note: The central atom is often written first, but not always. It is
usually the less electronegative element.
PCl3
CH2Cl2
HCN
BrO3 -