GENERAL CHEMISTRY
General, Organic, and Biochemistry
9th Edition
Katherine J. Denniston
Joseph J. Topping
Danaè R. Quirk Dorr
Robert L. Caret
Copyright © 2017 McGraw-Hill Education. Permission required for reproduction or display
3.1 Chemical Bonding
• Chemical bond - the force of attraction
between any two atoms in a compound
• This attractive force overcomes the
repulsion of the positively charged nuclei of
the two atoms participating in the bond
• Interactions involving valence electrons are
responsible for the chemical bond
3.1 Chemical Bonding
Lewis Symbols
1
• Lewis symbol - a way to represent atoms
using the element symbol and valence
electrons as dots
• As only valence electrons participate in
bonding, this makes it much easier to work
with the octet rule
• The number of dots used corresponds
directly to the number of valence electrons
located in the outermost shell of the atoms
of the element
3.1 Chemical Bonding
Writing Lewis Symbols
• The four sides around the atomic symbol
can each have two dots for a maximum of
eight (octet of electrons).
• Writing Lewis symbols
– Place one dot on each side until there are four dots
around the symbol
– Then add a second dot to each side in turn
– The number of valence electrons limits the number of
dots placed
– Each unpaired dot (unpaired valence electron) is
available to form a chemical bond
3.1 Chemical Bonding
Lewis Symbols for
Representative Elements
3.1 Chemical Bonding
Principal Types of Chemical Bonds:
Ionic and Covalent
2
• Ionic bond - attractive force due to the
transfer of one or more electrons from one
atom to another
• The attraction is due to the opposite charges of
the ions
• Covalent bond - attractive force due to the
sharing of electrons between atoms
• Some bonds have characteristics of both
types and not easily identified as one or the
other
3.1 Chemical Bonding
Ionic Bonding
2
• Representative elements form ions that
obey the octet rule
• Electrons are lost by a metal and they are
gained by a nonmetal
– Each atom achieves a “Noble Gas”
configuration
– 2 ions are formed; a cation and anion, which
are attracted to each other
• Ions of opposite charge attract each other
creating the ionic bond
3.1 Chemical Bonding
Ionic Bonding: NaCl
Consider the formation of NaCl
Na + Cl  NaCl
Sodium has a low
ionization energy (it
readily loses this
electron)
Na  Na+ + e−
When sodium loses the
electron, it gains the
Ne configuration
Chlorine has a high
electron affinity
When chlorine gains
an electron, it gains
the Ar configuration

..
..



: Cl  e  : Cl :
..
 .. 
3.1 Chemical Bonding
Essential Features of Ionic Bonding
• Metals tend to form cations because they have
low I.E. and low E.A.
• Nonmetals tend to form anions because they have
high I.E. and high E.A.
• Ions are formed by the transfer of electrons
• The oppositely charged ions formed are held
together by an electrostatic force
• Reactions between metals and nonmetals tend to
form ionic compounds
3.1 Chemical Bonding
Ion Arrangement in a Crystal
• As a sodium atom loses one electron, it becomes a
smaller sodium ion
• When a chlorine atom gains that electron, it
becomes a larger chloride ion
• Attraction of the Na cation with the Cl anion
forms NaCl ion pairs that aggregate into a crystal
3.1 Chemical Bonding
Covalent Bonding
Consider the formation of H2:
2
H + H  H2
• Each hydrogen has one electron in its valance
shell
• If it were an ionic bond it would look like this:
H   H   H  H :


• However, both hydrogen atoms have an equal
tendency to gain or lose electrons
• Electron transfer from one H to another
usually will not occur under normal conditions
3.1 Chemical Bonding
Covalent Bond
• Instead, each atom attains a noble gas
configuration by sharing electrons
H   H  H : H
Each hydrogen
atom now has two
electrons around it
and attained a He
configuration
The shared
electron
pair is called a
Covalent Bond
3.1 Chemical Bonding
Covalent Bonding in Hydrogen
3.1 Chemical Bonding
Features of Covalent Bonds
• Covalent bonds form between atoms with
similar tendencies to gain or lose electrons
• Compounds containing covalent bonds are
called covalent compounds or molecules
• The diatomic elements have completely
covalent bonds (totally equal sharing)
– H2, N2, O2, F2, Cl2, Br2, I2
.. ..
.. ..
: F   F :  : F : F :
.. ..
.. ..
Each fluorine is
surrounded by 8
electrons – Ne
configuration
3.1 Chemical Bonding
Examples of Covalent Bonding
..
..
2H    O  H : O : H
..
..
2e– from 2H
6e– from O
2e– for H
8e– for O
H
.
..
4H    C  H : C : H


H
4e– from 4H
4e– from C
2e– for H
8e– for C
3.1 Chemical Bonding
Polar Covalent Bonding and
Electronegativity
• The Polar Covalent Bond
– Ionic bonding involves the transfer of
electrons
– Covalent bonding involves the sharing of
electrons
– Polar covalent bonding - bonds made up
of unequally shared electron pairs
Polar Covalent Bond
somewhat positively charged
somewhat negatively charged
..
..
H  F: H : F:


• The electrons spend
more time with fluorine
• This sets up a polar
covalent bond
• A truly covalent bond
can only occur when
both atoms are identical
These two
electrons are
not shared equally
3.1 Chemical Bonding
Polar Covalent Bonding in HF
• Fluorine is electron
rich = • Hydrogen is electron
deficient = +
• This results in unequal
sharing of electrons in
the pairs = polar
covalent bonds
3.1 Chemical Bonding
Electronegativity
• Electronegativity - a measure of the
ability of an atom to attract electrons in a
chemical bond
• Elements with high electronegativity have
a greater ability to attract electrons than do
elements with low electronegativity
• Consider the covalent bond as competition
for electrons between 2 positive centers
– The difference in electronegativity determines
the extent of bond polarity
• The most electronegative elements are found in the upper
right corner of the periodic table
• The least electronegative elements are found in the lower
left corner of the periodic table
electronegativity increases
electronegativity increases
3.1 Chemical Bonding
Electronegativities
of Selected Elements
3.1 Chemical Bonding
Electronegativity Calculations
• The greater the difference in electronegativity
between two atoms, the greater the polarity of
their bond
• Which would be more polar, a H-F bond or H-Cl
bond?
• H-F … 4.0 - 2.1 = 1.9
• H-Cl … 3.0 - 2.1 = 0.9
• The HF bond is more polar than the HCl bond
3.2 Naming Compounds and
Writing Formulas of Compounds
• Nomenclature - the assignment of a correct
and unambiguous name to each and every
chemical compound
3
• Two naming systems:
– ionic compounds
– covalent compounds
Writing Formulas of Compounds
3.2 Naming Compounds and
Formulas of Compounds
• A formula is the representation of the
fundamental compound using chemical
symbols and numerical subscripts
– The formula identifies the number and type
of the various atoms that make up the
compound unit
– The number of like atoms in the unit is
shown by the use of a subscript
– Presence of only one atom is understood
when no subscript is present
Writing Formulas of Compounds
3.2 Naming Compounds and
Ionic Compounds
• Metals and nonmetals usually react to form
ionic compounds
• The metals are cations and the nonmetals
are anions
• The cations and anions arrange themselves
in a regular three-dimensional repeating
array called a crystal lattice
• Formula of an ionic compound is the
smallest whole-number ratio of ions in the
substance
Writing Formulas of Compounds
3.2 Naming Compounds and
Writing Formulas of Ionic Compounds
from the Identities of the Component Ions
• Determine the charge of each ion
3
– Metals have a charge equal to group number
– Nonmetals have a charge equal to the group
number minus eight
• Cations and anions must combine to give a
formula with a net charge of zero
• It must have the same number of positive
charges as negative charges
Writing Formulas of Compounds
3.2 Naming Compounds and
Predict Formulas
Predict the formula of the ionic compounds
formed from combining ions of the
following pairs of elements:
1. sodium and oxygen
2. lithium and bromine
3. aluminum and oxygen
4. barium and fluorine
Writing Names of Ionic Compounds
from the Formula of the Compound
• Name the cation followed by the name of
the anion
• A positive ion retains the name of the
element; change the anion suffix to -ide
Writing Formulas of Compounds
3.2 Naming Compounds and
Writing Names of Ionic Compounds
from the Formula of the Compound
• If the cation of an element has several
ions of different charges (as with
transition metals) use a Roman numeral
following the metal name
• Roman numerals give the charge of the metal
• Examples:
• FeCl3 is iron(III) chloride
• FeCl2 is iron(II) chloride
• CuO is copper(II) oxide
Writing Formulas of Compounds
3.2 Naming Compounds and
Common Nomenclature System
• Use -ic to indicate the higher of the
charges that ion might have
• Use -ous to indicate the lower of the
charges that ion might have
• Examples:
• FeCl2 is ferrous chloride
• FeCl3 is ferric chloride
Writing Formulas of Compounds
3.2 Naming Compounds and
Stock and Common Names for
Iron and Copper Ions
Writing Formulas of Compounds
3.2 Naming Compounds and
Common Monatomic
Cations and Anions
• Monatomic ions - ions consisting of a
single charged atom
Writing Formulas of Compounds
3.2 Naming Compounds and
Polyatomic Ions
• Polyatomic ions - ions composed of 2 or
more atoms bonded together with an
overall positive or negative charge
– Within the ion itself, the atoms are bonded
using covalent bonds
– The positive and negative ions will be
bonded to each other with ionic bonds
• Examples:
• NH4+ ammonium ion
• SO42- sulfate ion
Writing Formulas of Compounds
3.2 Naming Compounds and
Common Polyatomic Cations and Anions
Writing Formulas of Compounds
3.2 Naming Compounds and
Name These Compounds
1. NH4Cl
2. BaSO4
3. Fe(NO3)3
4. CuHCO3
5. Ca(OH)2
Writing Formulas of Compounds
3.2 Naming Compounds and
Writing Formulas of
Ionic Compounds from the
Name of the Compound
• Determine the charge of each ion
• Write the formula so that the
resulting compound is neutral
• Example:
Barium chloride:
Barium is 2+, Chloride is 1Formula is BaCl2
3
Writing Formulas of Compounds
3.2 Naming Compounds and
Determine the
Formulas from Names
Write the formula for the following ionic
compounds:
1.
sodium sulfate
2.
ammonium sulfide
3.
magnesium phosphate
4.
chromium(II) sulfate
Writing Formulas of Compounds
3.2 Naming Compounds and
Covalent Compounds
• Covalent compounds are typically
formed from nonmetals
• Molecules - compounds characterized
by covalent bonding
• Not a part of a massive three-dimensional
crystal structure
• Exist as discrete molecules in the solid, liquid,
and gas states
Writing Formulas of Compounds
3.2 Naming Compounds and
Naming Covalent Compounds
1. The names of the elements are written
in the order in which they appear in
4
the formula
2. A prefix indicates the number of each
kind of atom
Writing Formulas of Compounds
3.2 Naming Compounds and
Naming Covalent Compounds continued
3. If only one atom of a particular
element is present in the molecule,
the prefix mono- is usually omitted
from the first element
Example: CO is carbon monoxide
4. The stem of the name of the last
element is used with the suffix –ide
5. The final vowel in a prefix is often
dropped before a vowel in the stem
name
Writing Formulas of Compounds
3.2 Naming Compounds and
Name These
Covalent Compounds
1. SiO2
2. N2O5
3. CCl4
4. IF7
Writing Formulas of Compounds
3.2 Naming Compounds and
Writing Formulas of Covalent
Compounds
• Use the prefixes in the names to determine
the subscripts for the elements
4
• Examples:
• nitrogen trichloride
• diphosphorus pentoxide
NCl3
P2O5
• Some common names that are used:
–
–
–
–
H2O
NH3
C2H5OH
C6H12O6
water
ammonia
ethanol or ethyl alcohol
glucose
Writing Formulas of Compounds
3.2 Naming Compounds and
Provide Formulas for These
Covalent Compounds
1. nitrogen monoxide
2. dinitrogen tetroxide
3. diphosphorus pentoxide
4. nitrogen trifluoride
3.3 Properties of Ionic and Covalent
Compounds
• Physical State
– Ionic compounds are usually solids at room 5
temperature
– Covalent compounds can be solids, liquids, and
gases
• Melting and Boiling Points
– Melting point - the temperature at which a
solid is converted to a liquid
– Boiling point - the temperature at which a
liquid is converted to a gas
3.3 Properties of Ionic and
Covalent Compounds
Physical Properties
• Melting and Boiling Points
– Ionic compounds have much higher melting points
and boiling points than covalent compounds
– A large amount of energy is required to break the
electrostatic attractions between ions
– Ionic compounds typically melt at several hundred
degrees Celsius
• Structure of Compounds in the Solid State
– Ionic compounds are crystalline
– Covalent compounds are crystalline or amorphous –
having no regular structure
3.3 Properties of Ionic and
Covalent Compounds
Electrolytes and Nonelectrolytes
• Solutions of Ionic and Covalent
Compounds
– Ionic compounds often dissolve in water,
where they dissociate - form positive and
negative ions in solution
– Electrolytes - ions present in solution
allowing the solution to conduct electricity
– Covalent solids usually do not dissociate and
do not conduct electricity - nonelectrolytes
3.3 Properties of Ionic and
Covalent Compounds
Comparison of Ionic vs. Covalent
Compounds
Ionic
Covalent
Often
Composed of
Electrons
Metal + nonmetal 2 nonmetals
Transferred
Shared
Physical state
Solid / crystal
Dissociation
Yes, electrolytes
Any / crystal
OR amorphous
No,
nonelectrolytes
Low
Boiling/Melting High
3.4 Drawing Lewis Structures of
Molecules and Polyatomic Ions
Lewis Structure Guidelines
6
1. Use chemical symbols for the various elements to
write the skeletal structure of the compound
–
–
–
–
The least electronegative atom will be placed in the
central position
Hydrogen always occupies terminal positions
Halogens occupy terminal positions, except when
more electronegative elements are present
Carbon often forms chains of carbon-carbon
covalent bonds
3.4 Drawing Lewis
Structures of Molecules
Lewis Structure Guidelines
2. Determine the number of valence
electrons associated with each atom in
the compound
–
–
–
Combine these valence electrons to
determine the total number of valence
electrons in the compound
Polyatomic cations, subtract one electron for
every positive charge
Polyatomic anions, add one electron for
every negative charge
3.4 Drawing Lewis
Structures of Molecules
Lewis Structure Guidelines continued
3. Connect the central atom to each of the
surrounding with single bonds
4. Next, complete octets of all the atoms
bonded to the central atom
• Hydrogen needs only two electrons
• Electrons not involved in bonding are
represented as lone pairs
• After the terminal atoms have an octet,
provide the central atom with an octet if
valence electrons are still available
3.4 Drawing Lewis
Structures of Molecules
Lewis Structure Guidelines
5. If there are not enough valence electrons
to give the central atom an octet, move
lone pair electrons from terminal atoms
to form a new bond with the central
atom.
–
Continue to shift the electrons until all atoms
have an octet.
6. Recheck that all atoms have the octet rule
satisfied and that the total number of
valance electrons are used
3.4 Drawing Lewis
Structures of Molecules
Drawing Lewis Structures of
Covalent Compounds
6
Draw the Lewis structure of carbon dioxide, CO2
1. Draw a skeletal structure of the molecule
Arrange the atoms in their most probable order
C-O-O
and/or
O-C-O
Find the electronegativity of O=3.5 & C=2.5
Place the least electronegative atom as the central
atom, here carbon is the central atom
Result is the O-C-O structure from above
3.4 Drawing Lewis
Structures of Molecules
Drawing Lewis Structures of
Covalent Compounds
Draw the Lewis structure of carbon dioxide, CO2
2. Find the number of valence electrons for each
atom and the total for the compound
1 C atom  4 valence electrons = 4 e2 O atoms  6 valence electrons = 12 e-
16 e- total
3. Use electron pairs to connect the C to each O
with a single bond
O―C―O
4. Place electron pairs around the atoms
:O―C―O:
This satisfies the rule for the O atoms, but not for C
3.4 Drawing Lewis
Structures of Molecules
Drawing Lewis Structures of
Covalent Compounds - continued
Draw the Lewis structure of carbon dioxide, CO2
5.
Redistribute the electrons moving 2 e- from each
O, placing them between C―O
O=C=O
In this structure, the octet rule is satisfied
• This is the most probable structure
• Four electrons are between C and O
• These electrons are share in covalent bonds
• Four electrons in this arrangement signify a double
bond
6.
Recheck the electron distribution
• 8 electron pairs = 16 valence electrons, number
counted at start
• 8 electrons around each atom, octet rule satisfied
3.4 Drawing Lewis
Structures of Molecules
Lewis Structure of Polyatomic
Anions
6
Draw the Lewis structure of carbonate ion, CO321. Draw a skeletal structure of the molecule
Carbon is less electronegative than oxygen
•
•
This makes carbon the central atom
Skeletal structure and charge:
2. 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
3.4 Drawing Lewis
Structures of Molecules
Lewis Structure of Polyatomic
Anions - continued
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 oxygen, but not
for the carbon
Move a lone pair from one of the O atoms to form
another bond with C
6.
3.4 Drawing Lewis
Structures of Molecules
Lewis Structures Practice
Using the guidelines presented, write Lewis
structures for the following:
1. H2O
2. NH3
3. CO2
4. NH4+
5. N2
3.4 Drawing Lewis
Structures of Molecules
Lewis Structure, Stability,
Multiple Bonds, and Bond
Energies
7
• Single bond - one pair of electrons are
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
3.4 Drawing Lewis
Structures of Molecules
Bond Energy and Bond Length
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
3.4 Drawing Lewis
Structures of Molecules
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
• Who is right? All of you!
• In some cases it is possible to write more
than one Lewis structure that satisfies the
octet rule for a particular compound
O:
:O
:
C
: :
:O
: :
O:
:
C
:O:
: O:
: :
:O
:
:
: O:
: :
3.4 Drawing Lewis
Structures of Molecules
CO3 Resonance
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
3.4 Drawing Lewis
Structures of Molecules
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  2 valence electrons = 2 e2 H atoms  1 valence electrons = 2 etotal 4 e– Resulting Lewis structure:
H – Be – H
3.4 Drawing Lewis
Structures of Molecules
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
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
Distributing the electrons results in
this Lewis structure
:F:
: :: :
•
F
F P
F
F
:
1 P atom  5 valence electrons = 5 e5 F atoms  7 valence electrons = 35 e40 e- total
F
F:
:F P
F:
:F:
: :
3.4 Drawing Lewis
Structures of Molecules
Expanded Octet
:
3.4 Drawing Lewis
Structures of Molecules
Lewis Structures and Molecular
Geometry; VSEPR Theory
• Molecular shape plays a large part in
determining properties and shape
8
• 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.
3.4 Drawing Lewis
Structures of Molecules
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
3.4 Drawing Lewis
Structures of Molecules
A Stable Exception to the Octet
Rule
• Consider BeH2
– Only 4 electrons surround the beryllium atom
– These electrons in the bonds to the two atoms
have minimal repulsion when located on
opposite sides of the structure
– Linear structure having bond angles of 180°
3.4 Drawing Lewis
Structures of Molecules
Another Stable Exception to the
Octet Rule
• Consider BF3
– There are 3 bonded atoms around the central atom
– These bonded atoms have minimal repulsion when
placed in a plane, forming a triangle
• Trigonal planar structure with bond angles of
120°
3.4 Drawing Lewis
Structures of Molecules
Basic Electron Pair Repulsion of
a Full Octet
• Consider CH4
– There are 4 bonded atoms 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
3.4 Drawing Lewis
Structures of Molecules
Basic Electron Pair Repulsion of
a Full Octet with One Lone Pair
• Consider NH3
– There are three bonded atoms and one lone pair (four groups)
– 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°bond angles
3.4 Drawing Lewis
Structures of Molecules
Basic Electron Pair Repulsion of
a Full Octet with Two Lone Pairs
• Consider H2O
– There are two bonded atoms and two lone pair (four groups)
– 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
Copyright © 2017 McGraw-Hill Education. Permission required for reproduction or display
3.4 Drawing Lewis
Structures of Molecules
Basic Procedure to Determine
Molecular Shape
1. Write the Lewis structure
2. Count the number of bonded atoms and lone
pairs around the central atom
3. If no lone pairs are present, geometry is:
•
•
•
2 bonded atoms - linear
3 bonded atoms - trigonal planar
4 bonded atoms - tetrahedral
4. If there are lone pairs, look at the arrangement
of the atoms and name the geometry. Names
include:
•
•
Angular
Trigonal pyramid
3.4 Drawing Lewis
Structures of Molecules
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
3.4 Drawing Lewis
Structures of Molecules
Determine the Molecular
Geometry
• PCl3
• SO2
• PH3
• SiH4
3.4 Drawing Lewis
Structures of Molecules
Lewis Structures and Polarity
• Polar molecules when placed in an electric field
will align themselves in the field
9
• 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
3.4 Drawing Lewis
Structures of Molecules
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
3.4 Drawing Lewis
Structures of Molecules
Guidelines for Determining
Polarity
• Molecules that have no lone pair on the central
atom, and all terminal atoms are the same are
nonpolar
• Molecules with one lone pair on the central
atom are polar
• Molecules with more than one lone pair on the
central atom are usually polar
3.4 Drawing Lewis
Structures of Molecules
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
3.5 Properties Based on Molecular
Geometry and Intermolecular Forces
• Intramolecular forces – attractive forces
within molecules
– Example: 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
10
3.5 Properties Based on Molecular
Geometry & Intermolecular Forces
Solubility
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
3.5 Properties Based on Molecular
Geometry & Intermolecular Forces
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
3.5 Properties Based on Molecular
Geometry & Intermolecular Forces
Water and Ammonia: Hydrogen
Bonds
• The attractive forces, called hydrogen
bonds, pull ammonia into water,
distributing the ammonia molecules
throughout the water, forming a
homogeneous solution
3.5 Properties Based on Molecular
Geometry & Intermolecular Forces
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
3.5 Properties Based on Molecular
Geometry & Intermolecular Forces
Water and Oil
• 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
3.5 Properties Based on Molecular
Geometry & Intermolecular Forces
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 (m.p.) of a solid
– Higher boiling point (b.p.) of a liquid
3.5 Properties Based on Molecular
Geometry & Intermolecular Forces
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