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Chapter 3 - Continued
Structure and Properties of Ionic
and Covalent Compounds
Denniston
Topping
Caret
6th Edition
22 September 2011
Comparison of Ionic vs. Covalent
Compounds
Ionic
Covalent
Composed of
Metal + nonmetal 2 nonmetals
Electrons
Transferred
Shared
Physical state
Solid / crystal
Dissociation
Yes, electrolytes
Any / crystal
OR amorphous
No,
nonelectrolytes
Low
Boiling/Melting High
Drawing Lewis Structures of
Covalent Compounds
Draw the Lewis structure of carbon dioxide, CO2
Draw a skeletal structure of the molecule
1. Arrange the atoms in their most probable order
C-O-O
and/or
O-C-O
2. Find the electronegativity of O=3.5 & C=2.5
3. Place the least electronegative atom as the central
atom, here carbon is the central atom
4. Result is the O-C-O structure from above
Electronegativities of Selected Elements
Drawing Lewis Structures
5. Find the number of valence electrons for each atom
and the total for the compound
1 C atom x 4 valence electrons = 4 e2 O atoms x 6 valence electrons = 12 e16 e- total
: :
: :
6. Use electron pairs to connect the C to each O with a
single bond
O:C:O
7. Place electron pairs around the atoms
:O:C:O:
This satisfies the rule for the O atoms, but not for C
Drawing Lewis Structures of
Covalent Compounds
8.
Redistribute the electrons moving 2 e- from each O,
placing them between C:O
: :
: :
C::O::C
9.
In this structure, the octet rule is satisfied
• This is the most probable structure
• Four electrons are between C and O
• These electrons are shared in covalent bonds
• Four electrons in this arrangement signify a double
bond
10. Recheck the electron distribution
• 8 electron pairs = 16 valence electrons, number
counted at start
• 8 electrons around each atom, octet rule satisfied
Lewis Structures of
Polyatomic Ions
• Prepare Lewis structures of polyatomic
ions as for neutral compounds, except:
• The charge on the ion must be
accounted for when computing the
total number of valence electrons
Lewis Structure of
Polyatomic Cations
Draw the Lewis structure of ammonium ion, NH4+
Draw a skeletal structure of the molecule
1.
2.
Ammonium has this structure and charge:
The total number of valence electrons is determined by
subtracting one electron for each unit of positive
charge
1 N atom x 5 valence electrons = 5 e4 H atoms x 1 valence electron = 4 e- 1 electron for +1 charge
= -1 e8 e- total
3.
Distribute these 8 e- around the skeletal structure
Lewis Structure of Polyatomic
Anions
Draw the Lewis structure of carbonate ion, CO32-
Draw a skeletal structure of the molecule
1.
Carbon is less electronegative than oxygen
•
•
2.
This makes carbon the central atom
Skeletal structure and charge:
The total number of valence electrons is determined by
adding one electron for each unit of negative charge
1 C atom x 4 valence electrons = 4 e3 O atoms x 6 valence electron = 18 e+ 2 negative charges
= 2 e24 e- total
3.
Distribute these e- around the skeletal structure
Lewis Structure of Polyatomic
Anions
Draw the Lewis structure of carbonate ion, CO324.
Distributing the electrons around the central carbon
atom (4 bonds) and around the surrounding O atoms
attempting to satisfy the octet rule results in:
5.
This satisfies the octet rule for the 3 oxygens, but not
for the carbon
Move a lone pair from one of the O atoms to form
another bond with C
6.
Lewis Structure, Stability, Multiple
Bonds, and Bond Energies
• Single bond - one pair of electrons is
shared between two atoms
• Double bond - two pairs of electrons are
shared between two atoms
• Triple bond - three pairs of electrons are
shared between two atoms
• Very stable
H : H or H - H
.. ..
O :: O or O  O
 
.. ..
N  N or N  N
Bond energy - the amount of energy
required to break a bond holding two
atoms together
triple bond > double bond > single bond
Bond length - the distance separating the
nuclei of two adjacent atoms
single bond > double bond > triple bond
Lewis Structures and Resonance
• Write the Lewis structure of CO32• If you look around you, you will probably see
the double bond put in different places
• In some cases it is possible to write more than
one Lewis structure that satisfies the octet rule
for a particular compound
..
:O:
:
:O:
::
..
:O:
:
..
..
..
..
..
: O : C :: O  : O : C : O :  : O :: C : O :




• Experimental evidence shows all bonds are
the same length, meaning there is not really
any double bond in this ion
• None of theses three Lewis structures exist,
but the actual structure is an average or
hybrid of these three Lewis structures
• Resonance - two or more Lewis structures
that contribute to the real structure
Lewis Structures and Exceptions
to the Octet Rule
1. Incomplete octet - less then eight
electrons around an atom other than H
– Let’s look at BeH2
1 Be atom x 2 valence electrons = 2 e2 H atoms x 1 valence electrons = 2 etotal 4 e– Resulting Lewis structure:
H : Be : H or H – Be – H
Odd Electron
2. Odd electron - if there is an odd number of
valence electrons, it is not possible to give
every atom eight electrons
•
Let’s look at NO, nitric oxide
•
It is impossible to pair all electrons as the
compound contains an ODD number of valence
electrons
N - O
Expanded Octet
3.
Expanded octet - an element in the 3rd period or
below may have 10 and 12 electrons around it
•
Expanded octet is the most common exception
•
Consider the Lewis structure of PF5
•
Phosphorus is a third period element
1 P atom x 5 valence electrons = 5 e5 F atoms x 7 valence electrons = 35 e40 e- total
•
Distributing the electrons results in this Lewis structure
Lewis Structures and Molecular
Geometry: VSEPR Theory
• Molecular shape plays a large part in
determining properties and shape
• VSEPR theory - Valance Shell Electron Pair
Repulsion theory
• Used to predict the shape of the molecules
• All electrons around the central atom arrange
themselves so they can be as far away from
each other as possible – to minimize electronic
repulsion
VSEPR Theory
• In the covalent bond, bonding electrons
are localized around the nucleus
• The covalent bond is directional, having
a specific orientation in space between
the bonded atoms
• Ionic bonds have electrostatic forces
which have no specific orientation in
space
Molecular Bonding
• Bonding pair = two electrons shared
by 2 atoms
– H:O
• Nonbonding pair = two electrons
belonging to 1 atom, pair not shared
– N:
• Maximal separation of bonding pairs
A Stable Exception to the Octet
Rule
• Consider BeH2
– Only 4 electrons surround the beryllium atom
– These 2 electron pairs have minimal repulsion
when located on opposite sides of the structure
– Linear structure having bond angles of 180°
Another Stable Exception to the
Octet Rule
• Consider BF3
– There are 3 shared electron pairs around the central
atom
– These electron pairs have minimal repulsion when
placed in a plane, forming a triangle
– Trigonal planar structure with bond angles of 120°
Basic Electron Pair Repulsion of
a Full Octet
• Consider CH4
– There are 4 shared electron pairs around the central
Carbon
– Minimal electron repulsion when electrons are placed at
the four corners of a tetrahedron
– Each H-C-H bond angle is 109.5°
• Tetrahedron is the primary structure of a full octet
Basic Electron Pair Repulsion of
a Full Octet with One Lone Pair
Consider NH3
• There are 4 electron pairs around the central
Nitrogen
• 3 pairs are shared electron pairs
• 1 pair is a lone pair
– A lone pair is more electronegative with a greater electron repulsion
– The lone pair takes one of the corners of the tetrahedron without being
visible, distorting the arrangement of electron pairs
• Ammonia has a trigonal pyramidal structure with 107° angles
Basic Electron Pair Repulsion of
a Full Octet with Two Lone Pairs
Consider H2O
• There are 4 electron pairs around the central oxygen
• 2 pairs are shared electron pairs
• 2 pairs are lone pairs
– All 4 electron pairs are approximately tetrahedral to each other
– The lone pairs take two of the corners of the tetrahedron without being
visible, distorting the arrangement of electron pairs
• Water has a bent or angular structure with 104.5° bond angles
Predicting Geometric Shape Using
Electron Pairs
Basic Procedure to Determine
Molecular Shape
1. Write the Lewis structure
2. Count the number of shared electron pairs and
lone pairs around the central atom
3. If no lone pairs are present, shape is:
•
•
•
2 shared pairs - linear
3 shared pairs - trigonal planar
4 shared pairs - tetrahedral
4. Look at the arrangement and name the shape
•
•
•
•
•
Linear
Trigonal planar
Bent
Trigonal pyramid
Tetrahedral
More Complex Molecules
Consider dimethyl ether
• Has 2 different central atoms:
• oxygen
• carbon
– CH3 (methyl group) has tetrahedral geometry (like methane)
– Portion of the molecule linking the two methyl groups would bond
angles similar to water
Determine the Molecular
Geometry
• PCl3
• SO2
• PH3
• SiH4
Lewis Structures and Polarity
• A molecule is polar if its centers of positive and
negative charges do not coincide
• Polar molecules when placed in an electric field
will align themselves in the field
• Molecules that are polar behave as a dipole (having
two “poles” or ends)
• One end is positively charged, the other is negatively
charged
• Nonpolar molecules will not align themselves in
an electric field
Determining Polarity
To determine if a molecule is polar:
• Write the Lewis structure
• Draw the geometry
• Use the following symbol to denote the polarity
of each bond
Positive end of the
bond, the less
electronegative atom
Negative end of the bond,
more electronegative atom
attracts the electrons more
strongly towards it
Practice Determining Polarity
Determine whether the following bonds and
molecules are polar:
1. Si – Cl
1. O2
2.
H–C
2. HF
3.
C–C
3. CH4
4.
S – Cl
4. H2O
Properties Based on Electronic
Structure and Molecular Geometry
• Intramolecular forces – attractive forces
within molecules – chemical bonds
• Intermolecular forces – attractive forces
between molecules
• Intermolecular forces determine many
physical properties
– Intermolecular forces are a direct consequence
of the intramolecular forces in the molecules
Solubility and Intermolecular
Forces
Solubility - the maximum amount of solute
that dissolves in a given amount of
solvent at a specific temperature
• “Like dissolves like”
– Polar molecules are most soluble in polar
solvents
– Nonpolar molecules are most soluble in
nonpolar solvents
• Does ammonia, NH3, dissolve in water?
• Yes, both molecules are polar
Interaction of Water and
Ammonia
• The - end of ammonia, N, is attracted to the + end of
the water molecule, H
• The + end of ammonia, H, is attracted to the - end of
the water molecule, O
• The attractive forces, called hydrogen bonds, pull
ammonia into water, distributing the ammonia
molecules throughout the water, forming a
homogeneous solution
Interaction of Water and Oil
• What do you know about oil
and water?
– “They don’t mix”
• Why?
– Because water is polar and
oil is nonpolar
• Water molecules exert their
attractive forces on other water
molecules
• Oil remains insoluble and
floats on the surface of the
water as it is less dense
Boiling Points of Liquids
and Melting Points of Solids
• Energy is used to overcome the
intermolecular attractive forces in a
substance, driving the molecules into a
less associated phase
• The greater the intermolecular force, the
more energy is required leading to
– Higher melting point of a solid
– Higher boiling point of a liquid
Factors Influencing Boiling and
Melting Points
• Strength of the attractive force holding the
substance in its current physical state
• Molecular mass
• Larger molecules have higher m.p. and b.p. than
smaller molecules as it is more difficult to convert a
larger mass to another phase
• Polarity
• Polar molecules have higher m.p. and b.p. than
nonpolar molecules of similar molecular mass due to
their stronger attractive force
Melting and Boiling Points –
Selected Compounds by Bonding Type