Chemistry 20 – Molecules and Compounds Unit Bonding
In this unit we will learn how molecules and compounds form, as well as how to name
them. Some very basic terms and rules that we must keep in mind before proceeding:
1. When at atom has lost or gained an electron, it will take on a charge and
become an ion.
2. The outer most electrons for any atom are termed the valence electrons.
3. The Octect Rule states the atoms will try to lose or gain electrons to become
the same as the nearest noble gas.
4. If the outer most level is full, the reactivity will be very low, and will increase
the further the electron is from full.
5. Cations are positive ions formed by losing electrons.
6. Anions are negative ions formed by losing electrons.
7. To predict if the ionic charge on an element, we refer to the periodic table and
determine what will be the easiest way for the atom to get the same number as
its closest noble gas. Sodium (Na) has one valence electron, so do you think it
will lose 1, or will it gain 7?
Topic 1 – Bond Types
Matter
 anything that has mass and occupies space.
 the smallest portion of matter is the atom.
Molecules
 are pure substances which are clusters of atoms
 these clusters are electrically neutral.
There are two types of molecules:
1. Elements - molecules made up of identical atoms of two non-metals.
- there are seven diatomic molecules
(hint! One way to remember these molecules is that they are always the
“gens”. This means hydrogen, nitrogen, oxygen and the hydrogens.)
diatomic means that they are composed of two atoms from the same
element. Ex. N2 (g), O2(g) , F2 (g), Cl2 (g), Br2 (l), I2 (s)
- There are 3 polyatomic ions (exists in 3 or more atoms from the same
element). Ex. O3 (g)(ozone), P4 (s), S8 (s)
2.
Compounds - molecules that consists of atoms from
different elements
There are two types of Compounds
i. molecular compounds: are composed of two or more different
non-metal elements. Ex. H20 (l), CO2 (g), C2H5OH (g)
(ethanol)
ii. ionic compounds: composed of positive and negative ions
(remember that ions are atoms that have gained or lost
elections creating an electrical charge for the atom)
**Note** Ionic compounds can be created by the
combining of two simple ions (ie. Na+ with Cl- = NaCl) or
by the combining of a simple ion with a complex
polyatomic ion (i.e. Na+ with CH3COO- = Na(CH3COO)
The force that holds two atoms together is called a chemical bond. Chemical bonds may
form by the attraction between a positive nucleus and negative electrons or the attraction
between a positive ion and a negative ion. Previously you learned about atomic structure,
electron arrangement, and periodic properties of the elements. Many of these properties
are due to the number of valence electrons. These same electrons are involved in the
formation of chemical bonds.
Atoms tend to be stable with full outer energy levels. Eight electrons in the valence
energy level results in a stable atom. In seeking stability atoms form bonds between one
another in order to fill the octet rule, also called the rule of eight. The octet rule is useful
in determining the type of ions likely to form. Elements on the right sight on the periodic
table tend to gain electrons to acquire the noble gas configuration; therefore, these
elements form negative ions. And similarly elements on the left side of the table lose
electrons and form positive ions.
Lewis Dot Structures
There are three major types of chemical bonds:
Ionic Bonds, Covalent Bonds, and Metallic Bonds
Ionic Bonds
An Ionic bond is the bond that results from an electrostatic attraction between a positive
(metal) and negative ion (nonmetal). If ionic bonds occur between metals and oxygen,
oxides form. Most other ionic compounds are called salts. Ionic compounds contain a
metallic cation (positively charged ion) and a nonmetallic anion (negatively charged ion).
Ex. A sodium atom has one valence electron and the chlorine has seven valence
electrons. When the two elements combine the sodium atom transfers an electron to the
chlorine atom. The result is that both atoms now have full outer energy levels, resulting
in two stable ions .
 
Na  +  Cl 
 
 
Na+ Cl  
(table salt)
Cl-1
Na
Sodium Chloride Crystal
+1
Na+1
Cl-1
Cl-1 Na+1
Cl-1
Na+1
Na+1
Cl-1 Na+1
Cl-1
Na+1
Na+1
Cl-1
Cl-1 Na+1
NaCl Crystal Schematic
Covalent Bonds
Covalent bonds are chemical bonds that result from the sharing of valence electrons.
Covalent bonding generally occurs when the elements are relatively close to each other
on the periodic table. The majority of covalent bonds form between nonmetallic
elements; there are two types of covalent bonds: polar covalent bonds and non-polar
covalent bonds.
Polar Covalent Bonds
Electrons in covalent bonds can be shared unequally between unlike atoms to form polar
covalent bonds (i.e. H2O (l)). For example in polar covalent bonds of water, the electrons
shared by the atoms spend a greater amount of time, on the average, closer to the oxygen
nucleus than the hydrogen nucleus creating a partial positive and negative charge at each
end (pole) of the molecule.
Non-Polar Covalent Bonds
Cl-1
Cl-1 Na+1
Another type of covalent bonds exist when electrons are shared equally between identical
atoms to form non- polar covalent bonds (i.e. Cl2 (g)). For example the seven diatomic
molecules discussed earlier are all examples of non-polar covalent bonding.
Electron dot Diagram
Lewis Structural Diagram
Note: a shared pair of electrons can be represented by a line.
All covalent bonds can be classified as single, double, or triple bonds; depending on the
number of electrons shared. How would you show CO2 or N2?
Metallic Bonds:
A metallic bond is a bond that holds pure metals together through an attraction between
electrons and positive ions. The metallic bond falls somewhere between the ionic bonds
and covalent bonds. Although metal atoms have at least one valence electron, they do
not share these electrons with other atoms nor do they lose these electrons to form ions.
Instead, in this crowded condition, the outer energy levels of the metal atoms overlap.
This “electron sea model” proposes that all metal atoms in a metallic solid contribute to
their valence electrons to form a “sea” of electrons. These electrons are not held by any
specific atom and can move easily from one atom to the next. When the atom’s outer
electrons move freely throughout the solid, a metallic cation is formed. Each such ion is
bonded to all neighboring metal cations by the sea of valence electrons. The important
difference is that elections are free to move. It is this factor that is responsible for the
important metallic property of electrical conductivity.
Valence Shell Electron Pair Repulsion Theory
Electron dot structures and Lewis structural formulas fail to reflect the 3-D
shapes of molecules in reality these molecules exist in 3-D. The VSEPR theory
allows us to predict these shapes. The VSEPR theory states that because
electron pairs repel, molecules adjust their shapes so that these pairs are as far
apart as possible. Unshared pairs of electrons are also important when
predicting shape.
Please look at the Powerpoint labelled 3-D Shapes of Molecules.
Topic 2 – Bond Properties
How to determine whether a bond is Ionic, Covalent, or Metallic
The character and type of chemical bond can be predicted by using the electronegativity
difference of the elements that are bonded. Remember that electronegativity is defined as
a measure of an atoms attraction for electrons in a chemical bond. As learned in unit one,
the trend for electronegativity on the periodic table increases as you go across a period
and decreases as you go down the family (or group).
Electronegativity
difference between
the two atoms
 1.7
Type of bond
between atoms
< 1.7
Polar Covalent
0
Non-Polar
Covalent
0
Metallic
Ionic
Descriptions of the
electrons in the bond
Transfer of electrons
between metal and nonmetal to form ions
Electrons shared unequally
between unlike atoms
Electrons shared equally
between identical nonmetallic atoms
Electrons moving freely
between metallic (+ve) ions
Example
Li = 1.0
F = 4.0
Difference = 3.0
they form an ionic bond
Li+ FBa = 0.9
I = 2.5
Difference = 1.6
they form a polar covalent
bond.
O = 3.5
O = 3.5
Difference = 0
they form an
K = 0.82
K = 0.82
Difference = 0
they form a metallic bond
Electronegativity and Bond Character

For identical non-metallic atoms, which have an electronegativty difference of 0,
the electrons in the bond are equally shared between the two atoms and the bond
is considered non-polar covalent, this is a pure covalent bond.

Chemical bonds between bonds of difference atoms are never really completely
ionic or covalent, and the character of the bond depends on how strongly the
bonded atoms attract electrons.

Bonding often is not clearly ionic or covalent. As the difference in
electronegativity increases, the bond becomes more ionic in character. An
electronegativity difference of 1.70 is considered 50% covalent and 50% ionic.
Physical Properties of Specific Bond Types in Solid State
Ionic Compounds
Physical Property
Melting Point
Characteristics of solids
Particles making up solid
Conductivity
Description
High melting points
Hard and brittle
Positive and negative ions
Good conductors of electricity (electrolytes) in aqueous solutions
Bonding Structure
Solubility
Highly organized crystal lattice (a grid of positive and negative ions)
All dissolve into independent ions water (polar molecule) however
some dissolve better than others.
The strength of the ionic bond depends on the ion size and the ion
charge. Ionic bonds are stronger for:


Ions of higher charge.
Ions of smaller size (the attraction has to work over a smaller
volume of space when the ions are smaller. This makes the
attraction more effective than if the attraction has to work over a
larger volume of space, as is the case with a larger ion).
1.Which has the stronger bond and why, MgCl2 or NaCl?
NaCl has Na+1 and Cl-1 .
MgCl2 has ions with charges of Mg+2 and Cl-1 .
MgCl2 has an ion of a higher charge. This will produce a
stronger attractive force between the ions of unlike charge.
Thus MgCl2 has the stronger ionic bond.
2.Consider the ions represented in
Figure A and in Figure B at the
right. Which figure represents
Figure A
Figure B
ions that would have a stronger
ionic bond? Explain.
The ions in Figure A would give the stronger bond due to
their smaller size. The attraction between the unlike
charges would have to act over a smaller volume of space
making it more effective. The ionic bond would be stronger
for the smaller ions.
3.Which compound will have the stronger bond, NaCl or KCl?
Note the ion charges in both compounds are the same.
Thus, you need to compare ion sizes. The Cl- ion is the
same in both compounds so we really only have to
compare the size of the Na+ ion with the K+ ion. Look on
the periodic table. You will find Na above K. Since ions get
larger as you go down a family of elements in the periodic
table, you can safely say that K+ is the larger ion. Bonds
are stronger for ions of smaller size. Thus, Na+ has the
smaller size and therefore, NaCl will have a stronger bond
than KCl.
To draw a lewis structure or electron dot diagram for an ion you will
need to follow the following steps:
1. Find the element on the periodic table
2. Determine the nearest noble gas
3. Determine the number of electrons lost or gained so that the
element will have the same number of electrons as the nearest
noble gas. Remember that noble gases have an especially stable
electron population. All elements will tend to gain, lose or share
electrons to become like their nearest noble gas. This is called
the octet rule.
4. Determine the charge that the ion will form.
5. Draw the lewis structure.





Remember that positive ions contain no dots; negative ions
contain 8 dots.
Write the symbol of the element
Place the appropriate number of dots around the element
Enclose this in square brackets
Write the charge on the ion as a superscript outside the square
brackets.
Electron Dot Diagrams for Ionic Compounds
The steps to writing the electron dot diagram of an ionic compound
are:



place the electron dot diagrams for the ions side by side
alternate the positive and negative ions if possible
ensure that the sum of the charges on all the ions in the
compound is zero
Examples:
1. Write the electron dot diagram of NaCl.
2. Write the electron dot diagram for MgCl2.
3. Write the electron dot diagram for Al203
Assingment 1:Ionic Lewis Dot
Draw the lewis structure for the ions made from the following
elements:
1. O
2. P
3. Cl
4. Ca
5. K
6. Na
7. Al
8. N
9. F
10.
11.
12.
13.
14.
15.
Li
KBr
CaCl2
Na2O
MgO
MgF2
Physical properties of ionic compounds can be predicted by comparing
the strength of the ionic bond. Compounds having stronger ionic bonds
will have a higher melting point, a higher boiling point and a lower
solubility in water.
The stronger the bond, the more energy will be needed to break the
bond and thus the melting point will be higher.
The stronger the bond, the lower the solubility. The weaker the bond
the higher the solubility. Basically, if the bond is weaker, it would be
easier for that compound to break apart and dissolve.
The ionic bond is stronger for compounds that contain ions of higher
charge and smaller size. The higher the charge, the greater the force
of attraction holding the ions together. The smaller the ions, the
stronger the ionic bond. This is due to the fact that with smaller ions,
the attractive force has to act over a smaller volume of space. This
makes the attraction stronger than if the ions were larger and the
attractive force would have to act over a larger volume of space.
Assignment 2: Predicting physical properties of Ionic Bonds
Answer the following questions and give a reason for your answer.
1. Which element forms ions of higher charge, Na or Mg?
2. Which element forms ions of smaller size, Na or Mg?
3. Which compound
MgCl2? Explain.
4. Which compound
CaO? Explain.
5. Which compound
MgO? Explain.
6. Which compound
Explain.
7. Which compound
8. Which compound
MgO? Explain.
9. Which compound
or CaO? Explain.
10. Which compound
MgCl2? Explain.
would have the stronger ionic bond, NaCl or
would have the stronger ionic bond, MgO or
would have the stronger ionic bond, ZnS or
would have the stronger ionic bond, NaCl or KCl?
has the lower solubility, NaCl or KCl? Explain.
would have the higher solubility in water, ZnS or
would have the higher solubility in water, MgO
would have the lower solubility in water, NaCl or
Covalent Molecular Compounds
Physical Property
Description
Relatively low in comparison to ionic (i.e. sugar to salt)
Melting and boiling points
Soft – i.e. paraffin wax, grease
Characteristics of solids
molecules
Particles making up solid
Poor
Conductivity
Electrons tightly bound to atoms or shared by atoms in covalent
Bonding Structure
bonds. Molecules weakly attracted to each other and are easily
displaced
Some can dissolve in water but others will only dissolve in organic
Solubility
compounds. Many do not dissolve in water very well.
Example: Grease can be dissolved by benzene
Covalent Network Compounds
Physical Property
Description
Very high melting points because each atom is bound by strong
Melting and boiling points
covalent bonds. Many covalent bonds must be broken if the solid is
to be melted and a large amount of thermal energy is required for
this.
Covalent network substances are hard and brittle. If sufficient force
Characteristics of solids
is applied to a crystal, covalent bonds are broken as the lattice is
distorted. Shattering occurs rather than deformation of a shape.
atoms
Particles making up solid
Poor conductors because electrons are held either on the atoms or
Conductivity
Bonding Structure
Solubility
Metallic Compounds
Physical Property
Melting and boiling points
Characteristics of solids
Particles making up solid
Conductivity
Bonding Structure
Solubility
within covalent bonds. They cannot move through the lattice.
Atoms are strongly bound in the lattice, and are not easily displaced.
Some can dissolve in water but others will only dissolve in organic
compounds. Many do not dissolve in water very well.
Description
High melting and boiling points. Strong forces of attraction exist
between particles. A large amount of thermal energy is required to
overcome the strong electrical forces between the positive ions.
These forces operate throughout the lattice.
Metals are dense. Metals are malleable and ductile. The distortion
does not disrupt the metallic bonding. Metals are lustrous. The
presence of free electrons causes most metals to reflect light (nonmetals are transparent).
Positive ions in an electron cloud.
Good conductors of heat and electricity.
The particles present in metals are tightly packed in the lattice.
Some can dissolve in water but others will only dissolve in organic
compounds. Many do not dissolve in water very well.
Specific Bond Type Characteristics and their Everyday Application
Ionic Compounds
Common Name
Formula
Lye
NaOH (s)
Baking Soda
NaHCO3 (s)
Chemical Name
Sodium Hydroxide
Sodium Hydrogencarbonate
Table Salt
Milk of Magnesia
Cream of Tartar
Sodium Chloride
Magnesium Hydroxide
Potassium Hydrogentartrate
NaCL (s)
Mg(OH)2 (s)
KHC4H4O6
Covalent Molecular Compounds
Common Name
Formula
Sugar
C12H22O11 (s)
Alcohol
CH3CH2OH
(l)
Nail Polish
CH3COCH3
Remover
(l)
Natural Gas
CH4 (g)
Covalent Network Compounds
Chemical Name
Sucrose
Ethanol
Application
Unclogs drains
Raises bread and cakes by giving
off CO2 when heated
Adds salty taste in food
Works as an antacid and laxative
Mixes with baking soda to make
baked goods rise
Acetone
Application
Sweetener
Component of alcoholic
beverages
Solvent
Methane
Heating Fuel
Common Name
Silica
Quartz
Formula
SiO2 (s)
Chemical Name
Silicon Dioxide
Diamond
C
Diamaunde
Application
Cutting tools, ceramics, high
technology equipment i.e. space
shuttle
Jewel, industrial applications
(grinding, polishing, coating),
cutting tools.
Covalent Bonding
A covalent bond occurs when two adjacent nonmetallic atoms choose
to obtain their full valence level of electrons through electron sharing.
The electrons that are shared between the two atoms makes the bond.
This type of bond is called a covalent bond.
Lewis Structures of Covalent Compounds
The steps to draw a lewis structure of a covalent molecule are:
1
Write the symbol for the central atom in the molecule.
.
2
Join on the other atoms using dashes.
.
3
Calculate the number of bonds. The formula to use is:
.
or more simply, this formula can be expressed as:
4 Add dashes to your diagram to make the number of bonds
. calculated in step 3 equal to the number of dashes shown in the
diagram.
5
Add dots to obey the octet rule.
.
6
Do a final count that all atoms obeyed the octet rule.
.
For example
Draw the lewis structure for H2O.
Steps:
Work
1. Write the symbol for the
central atom in the molecule.
O
2. Join on the other atoms using
dashes.
H-O-H
3. Calculate the number of
bonds. Note: by the time you
are finished making the lewis
structure, each hydrogen will
have 2 valence electrons and
oxygen will have 8. This
makes a total of 12 valence
# of bonds = (12-8)/2
electrons needed. The current
number of valence electrons # of bonds = 2
before bonding is 1 on each
hydrogen and 6 on oxygen,
making a total of 8.
4. Add dashes to your diagram
to make the number of bonds
calculated in step 3 equal to
the number of dashes shown
in the diagram.
H-O-H
More dashes were not needed.
5. Add dots to obey the octet
rule.
6. Do a final count that all atoms obeyed the octet rule. See it!
Draw the lewis structure for SO2.
Steps:
Work
1. Write the symbol for the central
atom in the molecule.
S
2. Join on the other atoms using
dashes.
3. Calculate the number of bonds.
Note: by the time you are finished
making the lewis structure, each
atom will have 8 valence electrons
making a total of 24. Currently,
each oxygen atom has 6 valence
electrons and sulfur has 6 as well.
This means before bonding you
O-S-O
# of bonds = (24-18)/2
# of bonds = 3
have a total of 18 valence
electrons.
4. Add dashes to your diagram to
make the number of bonds
calculated in step 3 equal to the
number of dashes shown in the
diagram.
You could add dashes in one of two
places. You could show either:
5. Add dots to obey the octet rule.
The dots will be added accordingly.
Note that the added dots on the
double bonded oxygen atom are
placed on the corners of the
imaginary box. The dots on the
single bonded oxygen atom are
placed normally, along the sides of
the imaginary box. Only 2 dots are
needed on the S to make a
complete octet for sulfur. More than
one diagram is possible. These are
known as resonance structures.
6. Do a final count that all atoms obeyed the octet rule. See it here!
There are two ways that atoms will not be able to obey the octet rule.
These are:
1. Having too few valence electrons to ever obtain an octet.




Hydrogen, Beryllium and Boron have two few valence
electrons to ever obtain a full octet. See this.
Hydrogen can have at most 2 valence electrons after it shares
its electron with another atom.
Beryllium will have 4 valence electrons after it has finished
bonding.
Boron will have 6 valence electrons after it shares its valence
electrons with other atoms.
2. Expanding the octet to have 10, 12 or 14 valence electrons instead
of 8.

elements in periods 3, 4, 5, 6 and 7 can expand their octet to
have 10, 12, or 14 valence electrons.
Polar Covalent Bonds
Covalent bonds involve the sharing of electrons between atoms.
Consider the term "sharing". Does this mean equal sharing or unequal
sharing?
Think about sharing a Kit Kat chocolate bar with a friend. Does this
mean that both of you get 2 fingers from this chocolate bar or could
you have 3 fingers and give your friend only 1 finger? Would that be
sharing? Sure! Sharing does not mean that the sharing was done
equally! The sharing could be equal or it could be unequal and still be
called sharing.
So too with electrons. If the electron pair that is shared between two
atoms spends equal time with each atom, then the electron pair is
equally shared and we have a nonpolar covalent bond with no ionic
character.
If the electron pair that is shared between two atoms spends more
time with one atom than the other, then we have a polar covalent
bond.
In a polar covalent bond, the atom that has the electron pair more of
the time will be more negative than a neutral atom. It is said to be
slightly negative and is given the label of
(meaning slightly
negative). The atom that gets the electron pair less of the time will be
lacking a negative some of the time so it is given the label
(meaning slightly positive). These labels are placed above the
elements symbols when you draw a structural formula for a
compound.
To draw the structural formula of a compound, you perform the first
four steps in drawing a Lewis Structure. Simply do not do the last step
of adding the dots.
Is the bond between oxygen and hydrogen polar? If so, label the
polarity.
Steps:
Your work
1. Find oxygen on the periodic
table.
Oxygen is element 8.
2. Find hydrogen on the periodic Hydrogen is element 1.
table.
3. Determine if these elements
have the same or very close
electronegativities. Knowing
the trend in electronegativities
will help you. (Look at a chart
if in doubt.)
Oxygen (3.44) has a higher
electronegativity than hydrogen
(2.20). Knowing the trends in
electronegativities will be good
enough. If you forget the trend,
the actual values from a reference
table can be used.
4. If the electronegativities are
different and less than 1.7,
the bond will be polar
covalent.
The electronegativity difference
(3.44-2.20=1.24) indicates that
the bond is polar covalent.
5. Draw the bond using dashes.
6. Label the polarity. The atom
with the higher
electronegativity gets the
label.
Metallic Compounds
Common Name
Formula
Nickel
Ni
Chemical Name
Nickel
Silver
Ag
Silver
Sodium
Na
Sodium
Application
Wire production, coins, batteries,
pigments for paints and
ceramics, stainless steel, metal
coatings, computer equipment
Photography, electrical circuits,
dentistry, batteries, jewelry
Baking soda, table salt, yellow
highway lights, used in
production of paper, glass, and
soap.
The Strength of a Metallic Bond
The strength of the metallic bond depends upon:

the size of the metallic ion. The smaller the ion size, the
stronger the metallic bond. This is due to the fact that the
smaller the ion, the easier it will be for the electrons in the
electron sea to hold the ions together since they would have less
volume over which they would have to act. Do you remember

the trend in ion sizes? If not, refresh your memory on the trends
down a family and across a period!
the number of valence electrons. More valence electrons in the
electron sea make a stronger metallic bond.
Your job . . .
Use the factors that control metallic bond strength to predict which of
the following have the stronger metallic bond. Support your answer by
giving an explanation.
1.
2.
3.
4.
5.
Na or Mg
Li or K
Mg or Cu
Ca or Mg
K or Ca
Attraction Forces between Molecules
Until now we have considered the bonding between atoms within molecules
(intramolecular forces).
Now, lets consider what holds molecules together
(intermolecular bonding). Remember that ionic substances and metallic substances do
not form molecules; they form highly organized crystal lattices and electron “sea
clouds”, respectively. Therefore intermolecular bonding only occurs in substances that
form covalent bonds.
In a covalent compound, the intramolecular bonds are quite strong, but the attraction
between individual molecules is relatively weak. These weak forces of attraction are
known as intermolecular forces. Some intermolecular forces are stronger than others,
however, intermolecular forces are weaker than intramolecular forces. There are two
types of intermolecular bonding: van der Waals forces and Hydrogen Bonding.
1. Van der Waals Forces: These forces are the weakest attractions between
molecules. They are named after the Dutch chemist Johannes van der Waals
(1837-1923). Two major van der Waals forces are dispersion forces and
dipole interactions.
a. (London) Dispersion Forces:
These are the weakest of all
intermolecular attractions. Dispersion forces are sometimes called
London forces after the German-American physicist, Fritz London,
who first described them.
Intra =
within
Inter =
between
or among
Generally speaking, the strength of dispersion forces increases as the
number of electrons in a molecule increases. The Halogens are an
example of molecules whose major attraction for one another is caused
by dispersion forces. Remember that the Halogen Family contains 5
of the 7 diatomic molecules that utilize dispersion forces. Dispersion
forces exist between all particles, but they only play a significant role
when there are no stronger forces of attraction acting on particles.
Dispersion forces are dominant in identical non-polar molecules.
Fluorine and Chlorine, with relatively few electrons, are gases at STP
(standard temperature and pressure). The larger number of electrons
in Bromine generate larger dispersion forces, making it a liquid at
STP. Iodine, with still a larger number of electrons is a solid at STP.
Recall that oxygen molecules are non-polar because electrons are
evenly distributed between the equally electronegative oxygen atoms.
Under the right conditions, however, oxygen molecules can be
compressed to a liquid. For oxygen to be compressed there must be
some force of attraction (dispersion forces) between its molecules.
These forces result from temporary shifts in the density of electrons in
the electron clouds. Remember that these elections are in constant
motion. When two non-polar molecules are in close contact,
especially when they collide, the electron cloud of one molecule,
repels the electron cloud of the other molecule and for a moment the
electron density around each nucleus is greater in one region of each
cloud. Each molecule forms a temporary dipole. When temporary
dipoles are close together a weak dispersion force exists between the
oppositely charged regions of the dipole. Due to the temporary nature
of the dipole, dispersion forces are the weakest. (see figure below)
b. Dipole-Dipole: The force between polar molecules is stronger and is
called a dipole-dipole force.
Polar molecules contain permanent dipoles; that is, some regions of a
polar molecule are always partially negative and some regions of the
molecule are always partially positive. Attractions between these
oppositely charged regions of polar molecules are dipole-dipole forces.
Neighboring polar molecules orient themselves so that oppositely
regions line up. When hydrogen chloride gas molecules approach, the
partially positive hydrogen atom in one molecule is attracted to the
partially negative chlorine atom in another molecule.
(see diagram below)
2. Hydrogen Bonds: These bonds are the strongest of the intermolecular forces.
A hydrogen bond is simply a dipole-dipole attraction that occurs between
molecules containing a hydrogen atom bonded to a small, highly
electronegative atom with at least one lone electron pair. Remember that
electronegativity is defined as a measure of an atoms attraction for electrons in
a bond. Hydrogen bonds are extremely important in determining the
properties of water and biological molecules like carbohydrates and proteins.
For a hydrogen bond to form, hydrogen must be bonded to either a fluorine,
oxygen, or nitrogen atom. These atoms are electronegative enough to cause a
large partial positive charge on the hydrogen atom, yet small enough that their
lone pairs of electrons can come close to hydrogen atoms. Consider, for
example a water molecule. In a water molecule, the hydrogen atoms have a
large partial positive charge and the oxygen has a large partial negative
charge. When water molecules approach, a hydrogen atom on one molecule is
attracted to the oxygen atom on the other molecule as shown on the diagrams
below.
Effect of Intermolecular Bonds on Physical Properties
The effect of the intermolecular bond on the melting and boiling points of compounds
can be seen by considering the following list of compounds of hydrogen with the group
7A, or group 17, elements.
Compound
H2O
H2S
H2Se
H2Te
Molecular Mass
18 g/mol
34 g/mol
81 g/mol
130 g/mol
Melting Point C
0
-85.5
-60.4
-48.9
Boiling Point C
100
-60.7
-41.5
-2.2
Normally, as the molecular mass increases (and hence the number of electrons), the
melting points and boiling points will also increase. This trend is found to be followed
for the compounds H2S to H2Te but as we can see, not for H2O. In fact the hydrogen
bonding is so strong in water that the melting and boiling points are much higher than
would have been predicted.
In general, the greater the intermolecular bond strength between the molecules of a
compound, the higher the melting and boiling point of that compound. This is due to the
increased energy which much be supplied to break the strong intermolecular bonds
between the molecules to allow the compound to melt and eventually boil.
Relative Strengths of Bond Types
Any comparison of the relative magnitude of the different types of attractive forces
should take into account various effects such as: molecular size, number of electrons,
and molecular shape. For this reason, the following ranking of bond strengths is an
approximation only. The scale is a relative scale from 0 to 100. Metallic Bonding varies
over such a wide range that it cannot be classified in this way, even approximately. The
melting points of the metallic elements vary from negative 39 C for mercury to 3410 C
for Tungsten.
1
5
10
50
100
Assignment 3 Bonding Lab
Please get a copy of the lab in print form. Your task will be to use jujubes and
toothpicks to build models of only the small covalent molecules.
What do you need to do?
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Perform the experiment. Show your teacher your models. Your
teacher will mark your models.
Answer the questions on the lab.
Compare your responses with those of your lab partner.
Hand in your finished lab.