Molecular Compounds
Writing names and Formulas
Two Types of Compounds
 Molecular compounds
• Made of molecules.
• Made by joining nonmetal atoms
together into molecules.
Chemical Formulas
• Shows the kind and number of atoms
in the smallest piece of a substance.
• Molecular formula- number and kinds
of atoms in a molecule.
• CO2
• C6H12O6
Molecular compounds
• made of just nonmetals
• smallest piece is a molecule
• can’t be held together because of
opposite charges.
• can’t use charges to figure out how
many of each atom
Easier
• Ionic compounds use charges to
determine how many of each.
• Have to figure out charges.
• Have to figure out numbers.
• Molecular compounds name tells you
the number of atoms.
• Uses prefixes to tell you the number
Prefixes
1 mono2 di3 tri4 tetra5 penta-
6 hexa7 hepta8 octa9 nona10 deca-
Prefixes
To write the Chemical Formulas write:
Prefix
name Prefix name -ide
Prefixes
Exception 1: We don’t write mono- if there
is only one of the first element.
Example: CO2
Carbon Dioxide
CO We do write mono- if the second
element is only one.
Carbon monoxide
Prefixes
Exception 2: No double vowels when writing
(oa oo)
Example: C6O8
Name Hexacarbon octoxide
NOT Hexacarbon octaoxide
Name These
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N2O
NO2
Cl2O7
CBr4
CO2
BaCl2
Write formulas for these
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diphosphorus pentoxide
tetraiodide nonoxide
sulfur hexaflouride
nitrogen trioxide
Carbon tetrahydride
phosphorus trifluoride
aluminum chloride
Acids
Writing names and
Formulas
Acids
• Compounds that give off hydrogen
ions when dissolved in water.
• Must have H in them.
• will always be some H next to an
anion.
• The anion determines the name.
Naming acids
• If the anion attached to hydrogen is
ends in -ide, put the prefix hydroand change -ide to -ic acid
• HCl - hydrogen ion and chloride ion
• hydrochloric acid
• H2S hydrogen ion and sulfide ion
• hydrosulfuric acid
Naming Acids
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If the anion has oxygen in it
it ends in -ate of -ite
change the suffix -ate to -ic acid
HNO3 Hydrogen and nitrate ions
Nitric acid
change the suffix -ite to -ous acid
HNO2 Hydrogen and nitrite ions
Nitrous acid
Name these
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HF
H3P
H2SO4
H2SO3
HCN
H2CrO4
Writing Formulas
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Hydrogen will always be first
name will tell you the anion
make the charges cancel out.
Starts with hydro- no oxygen, -ide
no hydro, -ate comes from -ic, -ite
comes from -ous
Write formulas for these
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hydroiodic acid
acetic acid
carbonic acid
phosphorous acid
hydrobromic acid
Covalent bonding
How does H2 form?
• The nuclei repel
+
+
How does H2 form?
• The nuclei repel
• But they are attracted to electrons
• They share the electrons
+
+
Covalent bonds
• Nonmetals hold onto their valence
electrons.
• They can’t give away electrons to bond.
• Still want noble gas configuration.
• Get it by sharing valence electrons with
each other.
• By sharing both atoms get to count the
electrons toward noble gas configuration.
Covalent bonding
• Fluorine has seven valence electrons
F
Covalent bonding
• Fluorine has seven valence electrons
• A second atom also has seven
F
F
Covalent bonding
Fluorine has seven valence electrons
 A second atom also has seven
 By sharing electrons

F F
Covalent bonding
Fluorine has seven valence electrons
 A second atom also has seven
 By sharing electrons
 Both end with full orbitals

F F
8 Valence
electrons
Covalent bonding
Fluorine has seven valence electrons
 A second atom also has seven
 By sharing electrons
 Both end with full orbitals

8 Valence
electrons
F F
Single Covalent Bond
• A sharing of two valence electrons.
• Only nonmetals and Hydrogen.
• Different from an ionic bond
because they actually form
molecules by sharing electrons.
• Example of Single Covalent Bond
•F2
F F
Double and Triple Covalent
Bonds
• If two atoms share two pairs of
electrons or 4 ve- they are called
double covalent bonds.
• Example: O2 O O
• If two atoms share three pairs of
electrons or 6 ve- they are called
triple covalent bonds.
• Example: N2 N N
Structural Formulas
• A structural formula is a molecular
model that uses letters as symbols
and bonds to show the relative
position of atoms to one another
Drawing Lewis Structures
1. Determine the central atom, Always
the one least in number in the
chemical formula or the atom with
the lowest electronegativity.
2. Draw the skeleton structure
3. Add up all the valence electrons for
the molecule.
Drawing Lewis Structures
4. Divide the total number of available
electrons by 2. This gives you the
number of bonding pairs.
5. Place one bonding pair between the cental
atom and everyother atom in the skeleton
structure.
6. Subtract the number of bonding pairs
used from the number of available
electrons pairs.
7. Place the remaining pairs around the
terminal atoms until each atoms fulfills
the octet rule.
8. Any leftover electrons go on the central
atom.
Examples
N
H
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NH3
N - has 5 valence electrons
H - has 1 valence electrons
NH3 has 5+3(1) = 8
8/2= 4 available pairs
Examples
• Draw one bond between the N and
each H.
• Subtract the number of lines from
your available pairs. 4 available pairs –
3 lines = one pair left.
H
H N H
Examples
• This leaves one available pair but
Hydrogen only needs two electrons to
be stable.
• So the lone pair must go on the
Nitrogen.
H
H N H
Multiple Bonds
• Sometimes atoms share more than
one pair of valence electrons.
• A double bond is when atoms share
two pair (4) of electrons.
• A triple bond is when atoms share
three pair (6) of electrons.
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HCN
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
10/2= 5 available pairs.
Draw one bond between the central
atom and each terminal atom.
HC N
HCN
• Subtract 5 available pairs – 2 used
pairs and you have 3 pairs left.
• Carbon needs 2 more pairs and
nitrogen needs 3 more pairs.
HC N
HCN

When you don’t have enough pairs
left, you have to make double or
triple bonds.
HC N
HCN
H C N
To make double or triple bonds you
take pairs of electrons off the terminal
atoms and share them with the central
atom.
Polyatomic Ions
• Polyatomic ions are charged, but with
in the ion there is covalent bonding.
• You Draw the Lewis structure the
same as with other molecules, but
what ever the charge number is, tells
you how many e- to add or subtract
from the total number of valence efor the molecule.
Resonance
• When more than one dot diagram with
the same connections are possible.
• NO2• Which one is it?
• Does it go back and forth.
• It is a mixture of both, like a mule.
• NO3-
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
 CO

C O
Coordinate Covalent Bond
When one atom donates both
electrons in a covalent bond.
 Carbon monoxide
 CO

C
O
VSEPR
• Valence Shell Electron Pair Repulsion.
• Predicts three dimensional geometry of
molecules.
• Name tells you the theory.
• Valence shell - outside electrons.
• Electron Pair repulsion - electron pairs
try to get as far away as possible.
• Can determine the angles of bonds.
VSEPR
• Based on the number of pairs of
valence electrons both bonded and
unbonded.
• Unbonded pair are called lone pair.
• CH4 - draw the structural formula
VSEPR
H
H C H
H
• Single bonds fill
all atoms.
• There are 4 pairs
of electrons
pushing away.
• The furthest
they can get
away is 109.5º.
4 atoms bonded
• Basic shape is
tetrahedral.
• A pyramid with a
triangular base.
• Same shape for
everything with 4
pairs.
H
H
109.5º
C
H
H
3 bonded - 1 lone pair
• Still basic tetrahedral but you can’t
see the electron pair.
• Shape is called
trigonal pyramidal.
H N H H
H
N
H
H
<109.5º
2 bonded - 2 lone pair
• Still basic tetrahedral but you can’t
see the 2 lone pair.
• Shape is called
bent.
H O
H
O
H
H
<109.5º
3 atoms no lone pair
• The farthest you can the electron
pair apart is 120º
H
H
C
O
3 atoms no lone pair
• The farthest you can the electron
pair apart is 120º.
• Shape is flat and called
trigonal planar.
120º
H
H
H
C
O
H
C
O
2 atoms no lone pair
• With three atoms the farthest they
can get apart is 180º.
• Shape called linear.
180º
O C
O
Hybrid Orbitals
Combines bonding with
geometry
Hybridization
• The mixing of several atomic orbitals to
form the same number of hybrid orbitals.
• All the hybrid orbitals that form are the
same.
• sp3 -1 s and 3 p orbitals mix to form 4
sp3 orbitals.
• sp2 -1 s and 2 p orbitals mix to form 3
sp2 orbitals leaving 1 p orbital.
• sp -1 s and 1 p orbitals mix to form 4 sp
orbitals leaving 2 p orbitals.
Hybridization
• We blend the s and p orbitals of the
valence electrons and end up with the
tetrahedral geometry.
• We combine one s orbital and 3 p
orbitals.
• sp3 hybridization has tetrahedral
geometry.
3
sp
geometry
• This leads to
tetrahedral shape.
• Every molecule with a
total of 4 atoms and
lone pair is sp3
109.5º
hybridized.
• Gives us trigonal
pyramidal and bent
shapes also.
2
sp
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hybridization
when three things come off atom
trigonal planar
120º
one p bond
Contains a double bond
sp hybridization
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when two things come off
one s and one p hybridize
Linear
Contains a triple bons
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 - 0.5 Covalent nonpolar
• 0.5 - 1.0 Covalent moderately polar
• 1.0 -2.0 Covalent polar
• >2.0 Ionic
• Use a table of electronegativities
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+
H
d-
Cl
• The Cl pulls harder on the electrons
• The electrons spend more time near the
Cl
Polar Molecules
Molecules with ends
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?
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HF
H2O
NH3
CCl4
CO2
Intermolecular Forces
What holds molecules to
each other
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
• Dispersion forces
• Dipole Interactions
• depend on the number of electrons
• more electrons stronger forces
• Bigger molecules
Dipole interactions
• 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.
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+ dH F
Dipole Interactions
+
d
d+
d-
d
Hydrogen bonding
• Are the attractive force 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
Hydrogen bonding
H O
H