MOLECULAR SHAPES

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COVALENT BOND
FORMATION
• When one nonmetal shares one or
more electrons with an atom of
another nonmetal so both atoms end
up with eight valence electrons
• There is a mutual attraction of
different nuclei to the electron’s
orbitals.
Characteristics of covalent
compounds: MOLECULES
1.Composed of non-metals that are sharing
electrons.
2. Composed of molecules
3. There are two forces involved:
Intramolecular – strong covalent bond with in the
molecules
Intermolecular – weak bonds between the molecules
Van der Waals forces
4. Have low Melting Points (weak
Van derWaals forces are breaking)
Example: Water
Strong covalent bonds between H & O,
make up the water molecules
Intramolecular
• Every atom has full
energy levels
HO
H
H
O
Intermolecular
Weak attractive forces between different
water molecules.
H
Intermolecular Forces
• They are what make solid and liquid
molecular compounds possible.
• The weakest are called van der Waal’s
forces - there are two kinds
1. (London) Dispersion forces
2. Dipole Interactions
(London) Dispersion Forces
• Depend on the number of electrons
• More electrons stronger forces
• Bigger molecules more electrons
• Fluorine is a gas
• Bromine is a liquid
• Iodine is a solid
Dipole interactions
• Occur when polar molecules are attracted to
each other.
• Slightly stronger than dispersion forces.
• Opposites attract but not completely hooked
like+in ionic
solids.
+
-
d d
H F
d d
H F
Dipole Interactions
+
d
d+
d-
d
Hydrogen bonding
• Are the attractive forces caused by
hydrogen bonded to F, O, or N.
• F, O, and N are very electronegative so it is
a very strong dipole.
• The hydrogen partially share with the lone
pair in the molecule next to it.
• The strongest of the intermolecular forces.
Hydrogen Bonding
d+ dH O
+
Hd
Diatomic Molecules (HONClBrIF) –
Molecules composed of two
atoms of the same element.
** All Diatomic molecules are purely covalent or non-polar covalent
because: there is an equal sharing between the two atoms.
1s1 1s1
EX: H2
H
H:H
H
or
H-H Bond angle (180)
LINEAR SHAPE
Both atoms strive to fill their 1s orbital so both hydrogen's attract the
pair of bonding electrons equally
+ :
+
* s-s bonding is the only
bonding that is nondirectional
Electrons spend time here; mutual
attraction for the same electron pair.
Forms a NEW
Molecular orbital
+
+
Fluorine has seven valence electrons
 A second atom also has seven
 By sharing electrons
 Both end with full orbitals

8 Valence
electrons
FF
Directional Bonding: Atoms approach at 2z
Other possible bonding ( s-p, p-p, s-d or p-d)
8 Valence
electrons
Linear shape
(bond angle 180)
Single Covalent Bond
• H2 and F2 are both examples of single covalent bonding.
• A sharing of two valence electrons.
• One pair of electrons shared between two atoms
How many unshared pairs
of electrons does I2 have?
I I
6
unshared
pairs
(Unshared pairs are also called lone pairs or
nonbonding pairs)
Multiple Bonds
• Sometimes atoms share more than one pair
of valence electrons.
• A double bond is when atoms share two
pair of electrons. (4 electrons)
• A triple bond is when atoms share three pair
of electrons. (6 electrons)
:O::O:
How many pairs of unshared
electrons does O2 have?
:O=O :
Shape: linear
:
Bond angle: 180
or
:
Ex: O2
:
:
Double Covalent Bond
:O=O:
(4)
Triple Covalent Bond
:
Ex: N2
N
:N:::N:
:N=N:
Bond angle: 180
Shape: linear
How many pairs of unshared electrons does N2 have?
(2)
Coordinate Covalent Bond
• When one atom donates both electrons in a
covalent bond.
• Carbon monoxide
• CO
CO
Coordinate Covalent Bond
When one atom donates both electrons
in a covalent bond.
 Carbon monoxide
 CO

C O
Coordinate Covalent Bond
When one atom donates both electrons
in a covalent bond.
 Carbon monoxide
:C=O:
 CO

C O
Triple Bond
Oxygen donates a pair of electrons so both
atoms now have 8 valence electrons.
:S. .
:
.
:O.
:
:
Ex: SO2 (Sulfur dioxide)
. O:
.
: :
: :
:
Ex: SO2 (Sulfur dioxide)
:O S::O:
Double bond
:
:
:O S=O:
:
: :
:
:
:
Ex: SO2 (Sulfur dioxide)
:
:O S::O:
Double bond
Sulfur donates a pair of electrons to oxygen
so all 3 atoms have complete octets!
Can this molecule be drawn another
way and still be the same molecule?
YES
RESONANCE
2 or more valid electron dot formulas
that can be written for a molecule.
Ex: Ozone (O3)
:
:
:
.
.
.
:O. O:
. O:
.
Shares this pair of electrons
RESONANCE
:O::O:O:
:O:O::O:
:
:
: :
:
: :
:
:
: : :
• 2 or more valid electron dot formulas that
can be written for a molecule.
:O=O O:
or
: : :
:
:
: :
:
Ex: SO2 (Sulfur dioxide)
:
:O S::O:
:O::S: O:
or
Drawing Lewis Dot Structures
• Draw a skeleton structure putting the first
atom written in the center (except Hydrogen)
• Add up all the valence electrons.
• Count up the total number of electrons to
make all atoms have a stable octet.
• Subtract.
• Divide by 2
• Tells you how many bonds - draw them.
• Fill in the rest of the valence electrons to fill
atoms up.
Examples
N
H
• NH3
• N - has 5 valence electrons wants
8
• H - has 1 valence electrons wants
2
• NH3 has 5+3(1) = 8 Valence
• NH3 wants 8+3(2) = 14 Valence
• (14-8)/2= 3 bonds
• 4 atoms with 3 bonds
Examples
•
•
•
•
Draw in the bonds
3 bonds = 6 electrons (8-6=2)
All 8 valence electrons are accounted for
Everything is full
H
H NH
Examples
•
•
•
•
•
•
•
•
HCN C is central atom
N - has 5 valence electrons wants 8
C - has 4 valence electrons wants 8
H - has 1 valence electrons wants 2
HCN has 5+4+1 = 10
HCN wants 8+8+2 = 18
(18-10)/2= 4 bonds
3 atoms with 4 bonds -will require multiple bonds
- not to H
HCN
• Put in single bonds
• Need 2 more bonds
• Must go between C and N
HC N
HCN
Put in single bonds
 Need 2 more bonds
 Must go between C and N
 Uses 8 electrons (2 more to add)
10 – 8 = 2

HC N
HCN
Put in single bonds
 Need 2 more bonds
 Must go between C and N
 Uses 8 electrons - 2 more to add
 Must go on N to fill octet

HC N
Polar Bonds
• When the atoms in a bond are the same, the
electrons are shared equally.
• This is a nonpolar covalent bond.
• When two different atoms are connected,
the atoms may not be shared equally.
• This is a polar covalent bond.
• How do we measure how strong the atoms
pull on electrons?
Electronegativity
• A measure of how strongly the atoms attract
electrons in a bond.
• The bigger the electronegativity difference
the more polar the bond.
• 0.0 - 0.4 Covalent nonpolar
• 0.5 – 2.0 Covalent polar
• >2.0 Ionic
How
to
show
a
bond
is
polar
Isn’t a whole charge just a partial charge
•
 d+ means a partially positive
 d- means a partially negative
d+
(2.1)
H
d-
Cl
(3.0)
difference = .9 (polar covalent)
*The smaller the difference the more covalent the bond.
• The Cl pulls harder on the electrons
• The electrons spend more time near the Cl
Polar Molecules
• Molecules with a positive and a negative end
• Requires two things to be true
 The molecule must contain polar bonds
This can be determined from differences in
electronegativity.
Symmetry can not cancel out the effects of the
polar bonds.
Must determine geometry first.
Is it polar?
Bond (electronegativity)
• HF
Bond
Molecule
H=2.1
F=4.0 (1.9) v. polar
polar
• H2O H=2.1
O=3.5 (1.4) v. polar
polar
• NH3 H=2.1
N=3.0 (.9)
polar
polar
• CCl4 C=2.5
Cl=3.0 (.5)
polar
non-polar
• CO2
O=3.5 (1.0)
v. polar
non-polar
C=2.5
MOLECULAR
SHAPES
OF
COVALENT
COMPOUNDS
VSEPR THEORY
What Vsepr means
Since electrons do not like each other,
because of their negative charges, they
orient themselves as far apart as possible,
from each other.
Nonbonding electron pairs on the center
atom strongly repel the bonding pairs,
pushing the bonding pairs closer together
This leads to molecules having specific
shapes.
EXAMPLE:
Linear
BeF2
•Number of atoms = 3
•Number of Bonds = 2
•Number of Shared Pairs of Electrons = 2
•Bond Angle = 180°
** All molecules with 2 atoms are linear
Bent #1
EXAMPLE:
H2O
•Number of atoms = 3
•Number of Bonds = 2
•Number of Shared Pairs of Electrons = 2
•Number of Unshared Pairs of Electrons = 2
•Bond Angle = 105°
EXAMPLE:
Bent #2
O3
•Number of atoms = 3
•Number of Bonds = 2
•Number of Shared Pairs of Electrons = 2
•Number of Unshared Pairs of Electrons = 1
•Bond Angle = 105°
Trigonal Planar
EXAMPLE:
GaF3
•Number of atoms = 4
•Number of Bonds = 3
•Number of Shared Pairs of Electrons = 3
•Number of Unshared Pairs of Electrons = 0
•Bond Angle = 120°
EXAMPLE:
NH3
Pyramidal
•Number of atoms = 4
•Number of Bonds = 3
•Number of Shared Pairs of Electrons = 3
•Number of Unshared Pairs of Electrons = 1
•Bond Angle = 107°
EXAMPLE:
Tetrahedral
CH4
•Number of atoms = 5
•Number of Bonds = 4
•Number of Shared Pairs of Electrons = 4
•Number of Unshared Pairs of Electrons = 0
•Bond Angle = 109.5°
#
ATOMS
2
SHAPE
linear
linear
3
bent
3
Trigonal
4
planar
CENTRAL
ATOM
BOND
ANGLE
no
180
No unshared
pairs
180
Unshared
pairs
105
No unshared
pairs
120
pyramidal
4
Unshared
pairs
5
tetrahedral
No unshared
pairs
107
109.5
BOND
POLAR or NONPOLAR
MOLECULE
POLAR or NONPOLAR
Use electronegativity chart
Difference: 0-.4 non polar
.5 – 1.7 polar covalent
> 1.7 Ionic
COVALENT MOLECULES SUMMARY
Depends on
bond
Non- polar (all
terminal atoms
same)
polar
Non-polar (all
terminal atoms same)
polar
Non-polar (all
terminal atoms same)

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