Chapter 10
Molecular Structure:
Solids and Liquids
Homework
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Assigned Problems (odd numbers only)
“Questions and Problems”
(begins on page 292)
10.1 to 10.53
“Additional Questions and Problems” 10.65 to
10.91 (page 325)
“Challenge Questions”
(page 327)
10.93 to 10.99,
Representing Valence Electrons with Dots:
Electron-Dot Formulas
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The valence electrons are the most
important (chemically) electrons in the
atom
These electrons are situated in the
highest, occupied principal energy level
(n)
They are farthest from the nucleus , away
from the stable (filled) inner core of
electrons
Determine the chemical properties of an
element
Representing Valence Electrons with Dots:
Electron-Dot Formulas
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In the main-group elements, the valence
electrons are the in the outermost s- and
p- orbitals
The Roman numeral at the top of each
column represents the number of valence
electrons in each member of that group
The valence electrons for the main-group
elements can be represented by Lewis
Structures
Representing Valence Electrons with Dots:
Electron-Dot Formulas
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The Lewis (Dot) structure consists of an
elements symbol with one dot per each
valence electron placed about the
element’s symbol
The representative elements in the same
group have the same number of valence
electrons
The maximum number of valence
electrons for any element is eight
Representing Valence Electrons with Dots:
Electron-Dot Formulas
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Only the noble
gases have the
maximum
number of eight
electrons
The exception is
helium with only
two valence
electrons
Octet rule
Duet rule
Representing Valence Electrons with Dots:
Electron-Dot Formulas
Example: dot structures for
representative elements in period 2
The valence electrons are in the
outer s and p orbitals
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outer electron configuration
ns1
Li•
ns2
Be•
•
ns2np1
ns2np2
ns2np3
•
••
•B• •C•
•
•
•N•
•
ns2np4
••
•O:
•
ns2np5
••
:F:
•
ns2np6
••
:Ne:
••
Representing Valence Electrons with Dots:
Electron-Dot Formulas
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The force of attraction between two
types of atoms is a chemical bond
The ionic bond is the attraction
between positive and negative ions
Compounds held together by ionic
bonds are called ionic compounds
Each element forms an ion with a
stable noble gas configuration
Lewis Structures for Ionic Compounds:
Electrons Transferred
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During ionic bond formation, ions form
when atoms of two elements are present:
An element that can lose electrons and an
element that can gain electrons
The most stable of all outer electron
configurations is based on the chemical
properties of the noble gases
By chemical reactions and compound
formation, elements can attain a stable
outer electron configuration
Lewis Structures for Ionic Compounds:
Electrons Transferred
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Metals: Low ionization energy, form
ions (cations)
Nonmetals: High ionization energy,
accept electrons (anions)
The representative elements lose or
gain electrons to form ions with
outer electron configuration of the
nearest noble gas
Electron Configuration of Ions
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An ion: An atom that is electrically
charged from loss or gain of
electrons
Atoms are neutral due to the equal
number of protons and electrons
Loss or gain of an electron will leave
a net charge on the atom
Electron Configuration of Ions

Consider sodium
Na
1s22s22p63s1
Loss of 1 e-
Na+
1s22s22p6
Electron configuration
of neon
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Can attain the noble gas configuration
(neon) if it loses 1 electron to obtain Na+
Noble gas core remains but it does not
have the same chemical properties of
neon
Electron Configuration of Ions

Consider chlorine
Cl-
Cl
1s22s22p63s23p5
Gain of 1 e-
1s22s22p63s23p6
Electron configuration
of argon
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Can attain the noble gas configuration
(argon) if it gains 1 electron to obtain ClAttains same E configuration of the
nearest noble gas but it does not have the
same chemical properties of argon
Electron Configurations
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We know for Representative Elements:
• Group 1A metals form ions with +1 charge
• Group 2A metals form ions with +2 charge
• Group 7A nonmetals form ions with -1 charge
• Group 6A nonmetals form ions with -2 charge
• Group 8A nonmetals do not form ions, in fact they
are extremely unreactive
Transition Elements: All of the elements of the d area
of the periodic table
• Elements differ in the number of electrons in the d
subshell
• Octet rule does not apply here
• Loss of electrons does not lead to the noble gas
structure
Ions of the Metals of Groups IA,
IIA, IIIA
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Metals form
cations by losing
enough electrons
to get the same
electron
configuration as
the previous noble
gas
A stable noble gas
core attained in
each case
Atom
Atoms
Electron
Config
Ion
Ions
Electron
Config
Na
[Ne]3s1
Na+1
[Ne]
Mg
[Ne]3s2
Mg+2
[Ne]
Al
[Ne]3s23p1
Al+3
[Ne]
O
[He]2s22p4
O-2
[Ne]
F
[He]2s22p5
F-1
[Ne]
Ions of the Nonmetals of Group
VIA, VIIA
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Nonmetals form
anions by gaining
enough electrons
to get the same
electron
configuration as
the next noble gas
A stable noble gas
core attained in
each case
Atom
Atoms
Electron
Config
Ion
Ions
Electron
Config
Na
[Ne]3s1
Na+1
[Ne]
Mg
[Ne]3s2
Mg+2
[Ne]
Al
[Ne]3s23p1
Al+3
[Ne]
O
[He]2s22p4
O-2
[Ne]
F
[He]2s22p5
F-1
[Ne]
Lewis Structures for Ionic Compounds:
Electrons Transferred
••
Sodium Chloride
• Each sodium loses
one electron to
• Formed from
achieve an octet
sodium and
chlorine atoms
• Each chlorine atom
gains one electron
• An ionic bond
to achieve an octet
forms consisting
of a sodium ion
• Formula is NaCl
(+ charge) and a
chloride ion
••
•• (- charge)
+
Na• +•Cl
Na [ F ]
••
••
••
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••
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Lewis Structures for Ionic Compounds:
Electrons Transferred
Magnesium
• Each magnesium
Chloride
loses two electrons
to achieve an octet
• Formed from
magnesium and
• Each chlorine atom
two chlorines
gains one electron
to achieve an octet
• An ionic bond
forms consisting
• Formula is MgCl2
of a magnesium
ion (2+ charge)
and
two chloride ions
(- charge each)
Lewis Structures for Ionic Compounds:
Electrons Transferred
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Bonding involves ONLY valence electrons
Transfer of the s and p valence electrons
achieves stability of the nearest noble gas
Illustrates the sequence of atoms
Shows atoms and their valence electrons
• How they are distributed in a molecule
• Use a dot or X to represent an electron
Li•
Be•
•
Li•
•B•
•
•
•C•
Li1+
•
••
•N•
•
••
:F:
•
••
•O:
•
••
:F:
•• 1[:F:]
••
•
••
:Ne:
••
Lewis Structures for Ionic Compounds:
Electrons Transferred

Give the number of valence electrons for
Mg, N, and Br. Draw the Lewis dot
symbol for each of these elements
Mg: 2 valence electrons
Mg
N: 5 valence electrons
N
Br: 7 valence electrons
Br
Covalent Lewis Structures:
Electrons Shared
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Chemical bonding also occurs between two
nonmetals
Since nonmetals do not readily lose
electrons, when two nonmetals bind
together the electrons are shared
A covalent bond is a pair of electrons
shared by two atoms
It is the binding force that results from
two nuclei attracting the same shared
electrons
Covalent Lewis Structures:
Electrons Shared
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A covalent Lewis structure: A 2-D
representation of how atoms are
covalently bonded together
Each covalent bond is represented by a
pair of dots (bonding electrons)
Must also show all unshared pairs of
(nonbonding) electrons
All valence e- from every atom in a
molecule must be accounted for in the
form of bonds or nonbonding pairs
Covalent Lewis Structures:
Electrons Shared
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The atoms in covalently
bonded molecules often
have bonding and
nonbonding electrons
The number of covalent
bonds that an atom forms
is equal to the number of
electrons needed to form a
noble gas configuration
(octet)
The exception is hydrogen
which needs only two
electrons
Shared electrons can be
also represented by dashes
Formation of a Hydrogen Molecule
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The simplest covalent
bonding condition
Hydrogen has one 1s
electron
H atom requires one
additional electron to
obtain the stable
noble gas
configuration of
helium
Each H atom
contributes its one
electron
The electron pair
shared by the two
atoms, forming
diatomic hydrogen H2
Covalent Lewis Structure: Electrons Shared
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Duet Rule
• Hydrogen wants two electrons to attain the
noble gas configuration of helium
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Octet Rule
• All other main group elements want 8
electrons to achieve the noble gas
configuration
• Filled valence shell is achieved by
gaining/losing electrons or by sharing
electrons
Covalent Lewis Structures:
Double and Triple Bonds
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A single covalent bond is where two atoms
share one pair of valence electrons
Many molecules exist that need two or
three pairs of electrons to provide a
complete octet of electrons per atom
Multiple covalent bonds: Covalent bonds
where two pairs or three pairs of valence
electrons are shared between the same
two atoms
Covalent Lewis Structures:
Double and Triple Bonds
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Two identical nonmetal atoms (diatomic
molecules)
Each atom will share valence electrons with
the other
The shared pair of electrons allow each atom
to achieve a stable noble gas configuration
This configuration can be achieved by a
single, double, or triple shared pair of
electrons
●●
●●
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●●●
●●●
Covalent Lewis Structures:
Double and Triple Bonds
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Single Bond is a chemical bond where two
atoms share one pair of valence electrons
Double Bond is a chemical bond where two
atoms share two pairs of valence electrons
Triple Bond is a chemical bond where two
atoms share three pairs of valence electrons
F
F
O
O
N
N
Sharing Electrons Between Atoms
of Different Elements
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Two nonidentical
nonmetal atoms
The number of
covalent bonds an
atom forms will equal
the number of
electrons needed to
form a stable, noble
gas configuration
Hydrogen follows the
duet rule and forms a
stable, helium noble
gas configuration
Bonding Behavior of Elements
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Oxygen has 6 valence
electrons and 2 octet
vacancies
Can complete its octet
by forming two
covalent bonds
Nitrogen has 5
valence electrons and
3 octet vacancies
Can complete its octet
by forming three
covalent bonds
O
Group 6A
N
Group 5A
Bonding Behavior of Elements
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Carbon has 4 valence
electrons and 4 octet
vacancies
Can complete its octet
by forming four
covalent bonds
Fluorine has 7 valence
electrons and 1 octet
vacancy
Can complete its octet
by forming one
covalent bond
C
Group 4A
F
Group 7A
Bonding Behavior of Elements
C
N
Group 4A
Group 5A
O
F
Group 6A
Group 7A
Writing Lewis Structures for
Covalent Compounds
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A dot structure is a two-dimensional
representation of a molecule to show how
atoms are joined together by covalent
bonding
To write a dot structure:
A bond is shown as a pair of dots or a dash
Dot structures also show the location of
electron pairs not used in bonds
All valence electrons from every atom in the
molecule or PA ion must be accounted for
Writing Lewis Structures for
Covalent Compounds
1) Determine the arrangement of
atoms within a molecule
• If there are three or more atoms,
the central atom (usually) appears
only once in the formula
• Halogens are often terminal atoms
(at the ends) unless it is combined
with O as in oxyacids
• Hydrogen is always a terminal
atom
Writing Lewis Structures for
Covalent Compounds
2) Determine the total number of
valence electrons
• For main-group elements the group
numbers equal the number of
valence electrons for the element of
that group
• If it is an anion, add one electron to
the total for each negative charge
• If it is a cation, subtract one
electron from the total for each
positive charge
Writing Lewis Structures for
Covalent Compounds
3) Attach the central atom to each
bonded atom by a pair of electrons
• Subtract two electrons (from total
valence) for each single bond drawn
in the structure
Writing Lewis Structures for
Covalent Compounds
4) Distribute the remaining electrons
• Add electrons to each atom
bonded to the central atom until
each has eight electrons (octet
rule)
• Exception: hydrogen (duet rule)
• Any extra electrons should go to
the central atoms
Writing Lewis Structures for
Covalent Compounds
5) If the central atom does not fulfill
the octet rule, share one or more
lone pairs between a terminal atom
and the central atom (form multiple
covalent bonds)
• Double or triple bonds are formed
ONLY when one or both of the atoms
are C, N, O or S
Writing Lewis Structures for
Covalent Compounds
Example: Create the electron-dot formula for
Fluorine gas = F2
• Determine the arrangement of the atoms
F
●
●
F
F
F
• Determine the total number of valence electrons
for the dot structure (2 × 7 = 14)
• Begin with bonding electrons
• Distribute the remaining electrons among the
two atoms, giving octets to each fluorine atom,
then proceed to lone pairs on each atom
Writing Lewis Structures for
Covalent Compounds and Polyatomic Ions
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Write a Lewis structure for the
following:
• NH4+
• SO42• CO
• SCN-
Writing Lewis Structures for
Covalent Compounds: CO
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Determine the arrangement of the atoms
C: 4 electrons
O: 6 electrons
Total valence electrons: 10
If octets are not complete, form one or
more multiple covalent bonds
C O
C O
Writing Lewis Structures for
Polyatomic Ions: NH4+
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Determine the
arrangement of the atoms
Total number of valence eN: 5 valence electrons
H: 4×1 valence electron
Positive Charge: Subtract
one electron
Total valence electrons: 8
H
H N H
H
+
Writing Lewis Structures for
Polyatomic Ions: SO42-
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Determine the arrangement of the atoms
Total number of valence electrons
S: 6 valence electrons
O: 4 x 6 valence electrons
Negative Charge: Add two electrons
Total valence electrons: 32
O
O S O
O
2-
Writing Lewis Structures for
Polyatomic Ions: SCN-
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
Determine the arrangement of the atoms
Total number of valence electrons:
• S: 6 electrons
• C: 4 electrons
• N: 5 electrons
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
Charge: add one electron
Total electrons: 16
If octets are not complete, form one or more
multiple covalent bonds
S C N
S C N
-
Resonance: Equivalent Lewis Structures
for the Same Molecule
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Resonance occurs when no single dot
structure adequately describes bonding
in molecule
Resonance structures are two or more
dot structures for a molecule or ion that
has the same arrangement of atoms
Contain the same number of electrons
Differ only in the location of electrons
The true structure is the average of the
individual structures
Writing Resonance Structures
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Write a Lewis structure for nitrate ion and
include resonance structures
Determine the arrangement of the atoms
Total number of valence electrons
(3 ×6) + 5 + 1 = 24 valence elec

Use a double-headed arrow to connect the
structures
O
O N O
-
O
O N O
O
O N O
Writing Resonance Structures
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Write a Lewis structure for ozone and
include resonance structures (O3)
Determine the arrangement of the atoms
Total number of valence electrons
3 ×6 = 18 valence elec

The actual structure is called a “resonance
hybrid” of the two contributing structures
Resonance structures
Hybrid
Writing Resonance Structures
H
H
H
C C
C C
H C
C H
H C
C H
C C
C C
H
H
H
H
H
Predicting the Shapes of Molecules
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(Lewis) Dot structures describe the
distribution of valence electrons
among bonding pairs and
nonbonding pairs
They do not give info on the 3-D
shape of the molecule
The 3-D arrangement of atoms in a
molecule is determined by its
molecular geometry
Predicting the Shapes of Molecules

A Lewis structure for water
H

••
O
••
H
Are the atoms arranged in a straight line
or do they form a v-shape?
••
H
O
••
H
Predicting the Shapes of Molecules
VSEPR Theory
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
The 3-D shapes of molecules and PA
ions result from the orientation of
atoms about the central atom
VSEPR: Valence Shell Electron-Pair
Repulsion Theory focuses on the
bonding and nonbonding electrons in
the valence shell of the central atom
Predicting the Shapes of Molecules
VSEPR Theory
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The central atom’s electrons play
an important role in determining
molecular shape
Bonding and nonbonding pairs of
electrons have a natural electrostatic
repulsion that pushes them as far
apart from one another as possible
Predicting the Shapes of Molecules
VSEPR Theory
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
It is this repulsion of electron groups
(regions of negative charge) that
causes a molecule to have a certain
shape: molecular geometry
The central concept of VSEPR theory
is that the electron pairs in the
valence shell of atoms form an
arrangement in space that minimizes
the repulsion between the electron
pairs: electron geometry
Electron Geometries
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Two electron groups
• 2 atoms attached
• Shape: Linear
• Bond angle: 180°
Three electron groups
• 3 atoms attached
• Shape: Trigonal Planar
• Bond angle: 120°
Four electron groups
• 4 atoms attached
• Shape: Tetrahedral
• Bond angle 109.5°
Molecular Geometries
Linear
• 2 electron groups
• 2 bonding groups, 0 lone pairs
• 2 atoms on opposite sides of
the central atom
• 180° bond angles

180°
Trigonal Planar
• 3 electron groups
• 3 bonding groups, 0 one pairs
• 3 atoms form a triangle around
the central atom
• 120° bond angles
120°
Molecular Geometries

Tetrahedral
• 4 electron groups
• 4 bonding groups, 0 lone
pairs
• 4 surrounding atoms form a
tetrahedron around the
central atom
• 109.5° bond angles
109.5°

Predicting the Shapes of molecules
VSEPR Theory
Example: beryllium fluoride
• 2 electron groups and 2 bonding groups
• 0 lone pairs
• Beryllium does not follow the octet rule
• Electron geometry: linear
• Angle between electron groups: 180°
• Molecular geometry: linear
180°
F
Be
F
Shapes of molecules and ions
VSEPR Theory

Example: boron trifluoride
• 3 electron groups, 3 bonding groups, 0 lone
pairs
• Boron does not follow the octet rule
• Electron geometry: trigonal planar
• Angle between electron groups: 120°
• Molecular geometry: trigonal planar
F
B
120°
F
F
Shapes of molecules and ions
VSEPR Theory
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
Example: carbon tetrafluoride
4 electron groups, 4 bonding groups, 0 lone pairs
Electron geometry: tetrahedral
Angle between electron groups: 109.5°
Molecular geometry: tetrahedral
F
C
F
F
109.5°
F
VSEPR Theory: Central Atoms with
Bonding Pairs and Lone Pairs
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
Often the 3-D shape (molec. geometry) is
not described the same as the electron
geometry
The electron-pair geometry around a
central atom includes the spatial
positions of all bond pairs and lone pairs
Only the arrangement of atoms describes
the molecular geometry of the molecule,
not the lone pairs of electrons because
you can’t actually see lone pairs, you can
only see the atoms
VSEPR Theory: Central Atoms with Bonding
Pairs and Lone Pairs
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
Example: ammonia gas
4 electron groups, 3 bonding groups, 1
lone pair
3 atoms attached to the central atom
A tetrahedral electron geometry
Angle between electron groups: 109.5°
A trigonal pyramidal molecular geometry
H
H
N
H
VSEPR Theory: Central atoms with Bonding
Pairs and Lone Pairs
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Example: Water
4 electron groups, 2 bonding groups, 2
lone pairs
2 atoms attached to the central atom
A tetrahedral electron geometry
Angle between electron groups: 109.5°
A bent molecular geometry
O
H
H
Electronegativity and Polarity


Electronegativity is the ability of an atom
in a molecule to attract bonding
electrons towards itself
The higher the element’s
electronegativity, the greater its ability to
attract electrons
Electronegativity and Polarity

Electronegativity
• Increases across a period (left to right) on
the periodic table
• Decreases down group (top to bottom) on
the periodic table

In general:
• Nonmetals have higher electronegativity
values than metals and nonmetals tend to
gain electrons
• Metals have lower electronegativity values
and tend to lose electrons
• This occurs when an ionic bond is formed
Electronegativity

Fluorine (the reference element) is most electronegative,
Cesium (Francium) is least electronegative
Polarity of Bonds



A covalent bond involves pairs of
electrons shared equally between
two atoms
How they are shared (equally or
unequally) depends on the electron
donating and electron attracting
nature of the atoms
As the electronegativity difference
increases between two elements,
bond polarity increases
Polarity of Bonds

Nonpolar (Pure) covalent bond
• Two identical atoms will share the
bonding electrons equally

Polar covalent bond
• Two different atoms will not share the
bonding electrons equally
• One atom will have a greater attraction
for the shared pair than the other atom

Ionic bond
• Results from the transfer of one or more
electrons from one atom to another
Electronegativity



Bond polarity measures the amount
of unequal sharing of electrons in a
chemical bond
The difference in electronegativity
values determines the extent of
polarity in a bond
As the bond polarity increases, the
bond becomes more ionic
Electronegativity




General guidelines related to electronegativity
differences
Bonds with a difference of zero are nonpolar
(pure) covalent
Bonds with differences greater than 0.4 but less
than 2.0 are called polar covalent bonds
Bonds with differences 2.0 or greater are called
ionic
Molecule
O2
CO
Mg2O
Electronegativity
Values
3.5 and 3.5
2.5 and 3.5
1.2 and 3.5
Electronegativity
Difference
0.0
1.0
2.3
Bond Type
Pure Covalent
Polar Covalent Ionic
Electronegativity



Increases across period (left to right) on Periodic Table
Decreases down group (top to bottom) on Periodic Table
Larger difference in electronegativities means more polar
bond
• Negative end toward more electronegative atom
Molecule
O-O
C-O
Mg-O
Electronegativity
Values
3.5 and 3.5
2.5 and 3.5
1.2 and 3.5
Electronegativity
Difference
0.0
1.0
2.3
Bond Type
Pure Covalent
Polar Covalent Ionic
Polarity of Bonds/Dipole Moments



In a polar covalent bond: One of the
two different elements will inevitably
have a greater attraction for the
shared pair than the other
This unequal sharing causes the
entire molecule to behave like an
electric dipole (dipole moment)
Dipole: A body with two poles, one
partially negative and one positive
Polar Bonds and Polar Molecules





Molecules (as well as bonds) can have
polarity
A polar molecule has an unsymmetrical
distribution of electronic charge: The
bonding electrons are more attracted to
one part of the molecule more than other
parts
The polarity of a molecule has two factors:
Bond polarities
Molecular geometry
Polar Bonds and Polar Molecules


If a molecule contains only polar bonds
does not mean the molecule as a whole is
polar
Polarity will depend on whether or not the
polarity effects cancel each other

H
O 
Center of
positive
charge
Center of
negative
charge
H

(a)
Fig11_5
(b)
Polar Bonds and Polar Molecules



In a water molecule,
oxygen draws the
shared pair of
electrons closer to
oxygen and partially
withdrawn from
hydrogen
Since the dipoles do
not cancel, water is a
polar molecule
Since the dipoles do
cancel, SO3 is a
nonpolar molecule

In a sulfur trioxide
molecule, oxygen
draws the shared pair
of electrons closer and
partially withdrawn
from sulfur

H
O 
Center of
positive
charge
Center of
negative
charge
H

(a)
Fig11_5
(b)

end



Kinetic molecular theory of matter states
that the particles present in any phase of
matter are in constant, random motion:
thermal energy
The thermal (kinetic) energy is
temperature dependent
The particles interact together through
attractions and repulsions that creates
potential energy within the particles



Thermal (kinetic) and potential
energy in a chemical system
influence the chemical properties of a
system
Kinetic energy gives the particles
their motion and tends to move the
particles away from each other
Potential energy is an attractive force
which attracts the particles together


The relative influence of kinetic and
potential energy is the main
consideration when KM theory is
used to explain the general
properties of the gas, liquid, or solid
states of matter
The type of energy that dominates
will influence a substance’s physical
state



Molecules in a liquid do not have all
the same kinetic energy
These differences in energy are due
to collisions between the molecules
The molecules near the surface,
above the average KE, can escape
the intermolecular forces holding
them in the liquid phase
Interactions between Molecules



Intermolecular Forces are forces that
act between a molecule and another
molecule
They are electrostatic forces that are
similar to the intramolecular forces
involved in covalent bonding
Much weaker forces but strong
enough to influence behavior of
liquids
Properties of Liquids and Solids




The solid state is predominated by
potential energy rather than by
kinetic energy
Particles are in a fixed position by
strong electrostatic attractions but
vibrate due to kinetic energy
Definite volume and shape
High density: the particles are
located as close as possible
Properties of Liquids and Solids






In the liquid state is not dominated by potential
energy or by kinetic energy
Particles are in not in a fixed position due to
kinetic energy and can slide over each other
The potential energy (cohesive force) is strong
enough prevent total separation
Assumes the shape of the container it occupies
Definite volume and indefinite shape
High density: the particles are not widely
separated but located as close as possible
Properties of Liquids and Solids






In the gaseous state is dominated completely by
kinetic energy
Particles of a gas are independent of each other
and move in a totally random manner due to
their kinetic energy
The attractive forces between particles have been
overcome by kinetic energy which allows particles
to travel in all directions
Assumes both the volume and shape of the
container it occupies
Indefinite volume and indefinite shape
Low density: the particles are widely separated
and relatively few particles per unit volume
Properties of Liquids and Solids


Liquid and solid phases have many similar
characteristics but gases are very different
The average distance between the particles only
slightly different in the solid and liquid but vastly
different in the gaseous state
Surface Tension and Viscosity




In a liquid, surface tension arises when
molecules on the surface act like a thin
membrane that allows objects to move on
the surface liquid
Due to forces of attraction that bind them
across the surface and pull them into the
body of the liquid
This inward force causes the surface to
contract to as small a surface as possible
Unlike the molecules within the body of
the liquid that experience forces equally in
all directions
Surface Tension and Viscosity



Surface tension results from
an imbalance in forces of
attraction
The magnitude of the
surface tension of a liquid is
affected by the strength of
its intermolecular forces
Water has strong
intermolecular forces and
thus, very high surface
tension
Surface Tension and Viscosity




Since liquids are incompressible and can
flow, they assume the shape of their
container
Some liquids easily flow and others do not
(e.g. water vs. molasses)
Viscosity is a measure of a fluid’s
resistance to flow
The stronger the intermolecular forces,
the greater the resistance to flow
Attractive Forces in Compounds

Intermolecular Forces
• Ionic compounds: Ions (+ and -)
are held together by ionic bonds
• Neutral molecules: One or more of
these forces hold molecules
together in liquids and solids:
• Dipole-dipole
• Hydrogen-bonding
• Dispersion
Attractive Forces in Compounds



Dipole-dipole: Attractions between polar
molecules
The nonsymmetrical distribution of the charge
causes the molecules to line up
Positive end of one directed toward negative end
of other
Attractive Forces in Compounds



Hydrogen bonding: A special type of dipoledipole interaction
Occurs between molecules that have a H atom
bonded to F, O, or N
The partially positive H and a lone pair of
electrons on another N, O, or F atom
Attractive Forces in Compounds




Dispersion (London) forces: Short-lived dipoles caused by
uneven shifts in electron density
The uneven shift causes one end of the molecule to be slightly
positive and one end slightly negative
This induces the same electron shift in adjacent molecules
which causes an attractive force
Only intermolecular force possible in nonpolar substances
Matter and Changes of State


Matter: Anything
that has mass and
occupies space
It is separated into
three categories:
Solid, Liquid, and
Gas
Matter and Changes of State

A change in state is the most
common type of physical change
• Melting/Freezing
• Vaporization/Condensation
• Sublimation/Deposition

The composition of the substance
does not change, only its appearance
Melting/Freezing


A change of state requiring the input of heat
Heat of Fusion
• Heat energy required to melt 1 g of a
substance
• Heat energy that must be removed to
freeze 1 g of a substance
• Heat energy (to melt) 1 g of ice
(water at 0 °C):
heat of fusionwater  334 gJ
Change of State Problem


Calculate the heat
needed (in Joules) to
melt 15 g of ice at
0°C, and to heat the
water to 75 °C
Two parts to the
problem
• Melt ice (use heat
of fusion for water)
• Heat water (use
specific heat of
Water)
Change of State Problem


Melt the ice
Calculate the heat absorbed to melt the ice
at 0 °C (no change in temperature)
ΔH fus
J
 334
g
334 J
15 g 
 5010 J
g
Change of State Problem


Heat the water
Calculate the heat energy needed to warm the water from
0 °C to 75 °C
SH H2O
J
 4.184
g  C
SH H2O  mass (g)  Δt
4.184 J


 (75(75
C °C)
0 C)  4710 J
15.0 g 

g C
Change of State Problem
Combining Energy Calculations

Calculate the total heat
• Melting the ice (Q1)
• Heating the water (Q2)
Q1  Q2  Qtotal
Q1
5,010 J
Q2
4,710 J
Qtotal 9,720 J
Change of State

Sublimation: A phase change from solid to gas
without going through the liquid state
• Requires the absorption of heat
• No temperature change occurs during process
• Deposition is the reverse process (heat is released)

Evaporation: A phase change from a liquid to a
gas
• Requires the absorption of heat
• No temperature change occurs during process
• Condensation is the reverse process (heat is
released)
Change of State



Boiling: A special form of evaporation
where the liquid converts to vapor through
bubble formation
Boiling point: Temp at which the vapor
pressure of the liquid is the same as the
atmospheric pressure
This allows the bubbles at the liquid surface
to escape in the atmosphere
Vaporization/Condensation

Heat of Vaporization
• Heat energy required to vaporize
1 g of a substance
• Heat energy that must be removed
to condense 1 g of a substance
• Heat energy (to vaporize) 1 g of
water to vapor:
heat of vaporizati onwater  2260 gJ
Change of State Problem
Calculate the heat
needed (in Joules) to
heat 15 g of water
from 75 °C to 100 °C,
and to convert it to
steam at 100 °C
 Two parts to the
problem
• Heat the water (use
specific heat for
water)
• Convert water to
steam (use heat of
vaporization for
Water)
Change of State Problem


Heat the water
Calculate the heat energy needed to warm the water from
75 °C to 100 °C
SH H2O
J
 4.184
g  C
SH H2O  mass (g)  Δt
4.184 J



1570
15.0 g 

(
100
C

75
CJ
)
×25 °C

g C
Change of State Problem

Calculate the heat absorbed to convert the
liquid water to steam at 100 °C (no change
in temperature)
ΔHvap
J
 2260
g
J
15.0 g  2260  33,900 J
g
Change of State Problem
Combining Energy Calculations

Calculate the total heat
• Heat the water (Q1)
• Convert liquid water to steam (Q2)
Q1  Q2  Qtotal
Q1
1,570 J
Q2
33,900 J
Qtotal
35,500 J
Heats of Fusion and Vaporization
Heating Curve

Illustrates the steps involved in changing
a solid to a gas
• Heat added is shown on the x-axis
• Temperature is shown on the y-axis

Energy required to undergo a series of
phase changes depends on the (three)
property values of the substance:
• Specific Heat
• Heat of Fusion
• Heat of Vaporization
Heating Curve
No temp change
No temp
change

end
Sharing Electrons Between Atoms
of Different Elements



Two nonidentical
nonmetal atoms
The number of
covalent bonds an
atom forms will
equal the number
of electrons needed
to form a noble gas
configuration
Each vacancy +
unpaired electron
combination in the
valence shell can be
used to form a twoelectron bond

Each atom will share
valence electrons with
the other forming a
shared pair of bonding
electrons (achieves a
stable noble gas
configuration)
Covalent Lewis Structures:
Double and Triple Bonds Bonding



Single Bond
• Uses a single pair of electrons between two
atoms
Double Bond
• Uses two pairs of electrons between the
same two atoms
Triple Bond
• Uses three pairs of electrons between the
same two atoms
F
F
O
O
N
N
Writing Resonance Structures


Example: Write a Lewis structure for nitrate
ion and include resonance structures
Use a double-headed arrow to connect the
structures
-
O
O N O
O
O N O
O
O N O
Lewis Structures for Ionic Compounds:
Electrons Transferred


The reaction of sodium and chlorine
The reaction represented through dot
structures:
••
Na• + :Cl:
•
••
Li1+[:F:]1••
Resonance



Resonance occurs whenever it is
possible to draw two or more electrondot structures
They differ only in the location of a
double bond between the same two
types of atoms
Two of the four electrons in the double
bond rapidly move between the two
single bonds: The true structure is the
average of the individual structures