Intensified Chemistry
Unit I Notes: Chemical Bonding and Molecular Structure
Page 1 of 13
I-1 Reading
A chemical bond is a force of attraction between different atoms that results from a shift
in the location of valence electrons. The text describes this electron shift as a change in
electron arrangements. This bond forms due to a drive toward electronic stability that
we have already talked about—the octet rule. The more stable the electronic
configuration, the lower will be the energy of a system. We have already discussed how
atoms with a full valence shell are the most stable, so valence electrons will shift in such
a way that there are eight electrons in the valence shell.
Ionic bonds can be defined as the attractive force that forms when the valence shell
electrons are transferred from the valence shell of one atom to the valence shell(s) of
another atom(s). The atoms that lose electrons become cations, and those that gain
electrons become anions. The positive and negative charges of these species hold them
together because of Coulombic forces of attraction. Because of these strong forces of
attraction, ionic compounds are characterized by high melting points and the ability to
conduct electricity in the molten (high temperature liquid) state. They also tend to be
soluble in water, and often form sharply-defined crystals in the solid state.
A covalent bond forms due to the sharing of valence electron pairs between the nuclei
of two atoms—the positively charged nuclei are each attracted to the concentrated
negative charge of the shared pair of electrons that is between them. This attraction is
what bonds the atoms together. Once again, the sharing of electrons allows each atom to
have 8 electrons in its valence shell, thus lowering the energy state of the contributing
atoms. The hypothetical line between the nuclei in a covalent bond is defined as the bond
axis. When an atom is bound to more than one other atom, the angle between the two
bond axes is defined as the bond angle. The distance between the nuclei along the bond
axis is called the bond length. Because the shared electrons are not exactly stationary,
the bond can be thought of as a spring that rotates, bends, and stretches as if it were
vibrating about the most probably electron pair locations. This means that bond lengths
and angles are average values.
We use electronegativity (EN) values to determine what type of bond will form. There
are two different calculations that we can use: Percent ionic character (% IC) or
difference in electronegativity (EN) values:
%IC 
HIGH _ EN  LOW _ EN
HIGH _ EN
100%
EN = HIGH EN  LOW EN
When EN > 1.67, (this corresponds to %IC > 50%), the bond is characterized as ionic.
When EN < 1.67, (this corresponds to %IC ≤ 50%), the bond is characterized as
covalent.
The covalent bond category can be further broken down into pure covalent bonds and
polar covalent bonds. When the electron pair(s) are equally shared, they will be equally
attracted to both nuclei in the bond. This means that the electronegativity values for each
of the atoms involved in the bond be equal or nearly equal. It is clear that if the EN values
are the same (such as for the diatomic molecules that we discussed in the last unit), the
%IC will be 0%. In fact, we generally classify any bond with %IC ≤ 5% (EN < 0.3) as
being nonpolar, or pure covalent bonds. When 5% ≤ %IC ≤ 50% (0.3 ≤ EN ≤ 1.67),
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Unit I Notes: Chemical Bonding and Molecular Structure
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the shared pair of electrons will be closer to the nucleus with the higher EN, and that side
of the bond has more negative charge than the nucleus with the lower EN. This type of
bond is called a polar covalent bond. We designate the atom with the higher EN as
being slightly negative with a δ symbol, and the atom with the lower EN as being
slightly positive with a δ+ symbol. The following is a chart that organizes the above info:
%IC Range
EN Range
Bond Classification
%IC ≤ 5%
EN < 0.3
nonpolar or pure covalent
5% ≤ %IC ≤ 50%
0.3 ≤ EN ≤ 1.67
polar covalent
%IC > 50%
EN > 1.67
ionic
%IC is generally considered the more reliable measure to determine what kind of bond
will form, so it is what we will rely on for our calculations.
Now that we have defined the bond types, we can be more specific about what a
molecule is—it is a neutral group of atoms that are held together with covalent bonds.
This means that chemical compounds whose simplest units are molecules are called
molecular compounds. We already know that a molecular formula indicates the
relative numbers of atoms of each element in a single molecule. A diatomic molecule is
formally defined as a molecule with only two atoms, but we usually use that term to refer
to the elemental form of the seven elements that we discussed in the last unit: N2, O2, F2,
Cl2, Br2, I2, and H2.
We already know a good bit about ionic compounds—they are composed of positive and
negative ions that are combined so that the numbers of positive and negative charges are
equal. We also know that a single formula unit of an ionic compound is represented by
its empirical formula, or simplest whole number ratio of atoms possible. We do this
because ionic compounds are actually 3-D crystals of ions attracted to each other because
of their opposite charges. This 3-D structure can potentially go on and on forever, which
is very different from discrete molecules of covalent compounds.
We have already discussed in class how ionic compounds form, but let us review: metal
atoms are likely to give up electrons in order to obtain the noble gas electron
configuration of the previous noble gas element on the periodic table. This means that
sodium will give up 1 electron in order to gain the electronic structure of neon. In fact, all
group 1 elements will give up 1 electron to have a +1 charge. All group 2 elements will
give up 2 electrons to have a +2 charge. Simultaneously, nonmetal elements will be
looking to gain electrons. Fluorine will gain 1 electron to gain the electronic
configuration of neon, and achieve a charge of 1. You will notice that the octet rule is
also the driving force behind ionic bonding, but the large differences in
electronegativities lead the metals to completely give up their electrons, and the
nonmetals to acquire them.
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Unit I Notes: Chemical Bonding and Molecular Structure
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We can represent the formation of these ions with dot structures. Note the use of
brackets in each case. They are not used in the text, but I want you to use them for
clarity and understanding:
Na
[Na]+1 (notice that the electron dot is gone)
F
[: F :]1 (notice that the electron from Na is now
attached here)
We have already talked about how nature is always trying to minimize potential energy.
In ionic crystals, ions minimize their potential by packing themselves very closely
together into an orderly arrangement called a crystal lattice. The location of the ions with
respect to each other are determined by Coulombic attractive and repulsive forces. The
attractive forces arise between oppositely charged species, and the repulsive ones arise
between like charges. When the attractive and repulsive forces are exactly balanced, the
distance between ions in a lattice will be determined. This minimization of energy when a
crystal lattice is formed is characterized by its lattice energy, or the energy released (is
this an exothermic or an endothermic process?) when one mole of an ionic crystalline
compound is formed from gaseous ions. These values would therefore have the units of
kJ
.
mol
Metals behave very differently when they are bound together in so-called metallic
bonds. Metals are excellent conductors of thermal energy and electricity because their
electrons are very loosely held, and therefore are very mobile (they can move around
very easily). In most metals, the p-orbitals and d-orbitals are either empty or
incompletely occupied, and overlap with the orbitals of other metal atoms. This is what
allows electrons to move freely within a bulk metal. We say the electrons are delocalized
because they are not strictly associated with any given nucleus, and that these mobile
electrons form a sea of electrons around the metal atoms which are packed together in a
crystal. The chemical bonding resulting from the attraction between the atoms (read this
as “nuclei”) and the surrounding sea of electrons is called metallic bonding.
We now know how the sea of electrons accounts for the high thermal and electrical
conductivity of metals. There are also many orbitals separated by very small energy
differences, so metals can absorb across a wide range of light frequencies. When the
electrons inevitably lose this energy, they give light back off. This emission of small
amounts of energy as light is responsible for the shiny appearance of metals. Metal
ductility and malleability are a result of the fact that metallic bonding forces are
generalized—they are the same in all directions. This permits whole planes of metal
atoms to slide past one another fairly easily. This is clearly very different from what
happens with ionic crystals.
We characterize metallic bond strength with the heat of vaporization. This is the amount
of energy required to convert the solid state to individual atoms in the gaseous state. The
larger the heat of vaporization for a metal, the stronger are its metallic bonds.
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Unit I Notes: Chemical Bonding and Molecular Structure
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I-2 Readng
The octet rule is defined as: Chemical compounds tend to form so that each atom, by
gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied
energy level (the valence shell). The shared electron pair effectively fills each atoms
outermost energy level in order to achieve the most stable electron configuration.
There are a few exceptions to the octet rule. These occur in compounds of beryllium and
boron, which are both capable of forming covalent bonds. Because there are only two
electrons in the valence shell of beryllium, it will form only 2 covalent bonds, for a total
of 4 electrons in the outermost energy level. Similarly, boron will tend to form only 3
covalent bonds for a total of 6 electrons in the outermost energy level. There are also
exceptions to this rule when d-orbitals get involved (Group 15, and group 17 elements),
and we will discuss these exceptions later.
We have already discussed electron dot notation, and we use it extensively in this unit to
represent the different types of bonds that can form. Remember that Lewis electron dot
structures visualize the valence shell of an atom. There will therefore be between 1 to 8
dots surrounding the element symbol, with each dot representing an electron.
We now proceed to the use of dot structures to represent covalent bonds. Let us start with
H2, since it is the simplest case. There will be parts of this that are hand-drawn, so please
excuse this:
H : H
Notice the loop that I have placed around the two electron dots to indicate that they are
shared pairs. Now let us try a dot structure for fluorine gas, F2:
:F : F:
You will notice that only the shared pair of electrons has a loop around it. The three
remaining pairs of electrons around each of the fluorine symbol indicate unshared pairs
of electrons. A single shared pair of electrons between two atoms is called a single
covalent bond. We could also represent the F2 molecule with a structural formula as:
:F―F:
where the single dash also represents a shared electron pair between the fluorine atoms.
It is also possible to have more than one pair of electrons shared between two atoms. A
double covalent bond, or double bond is made when two pairs of electrons are shared
between two atoms. An example of this is given for the molecule ethane, C2H6 in the text.
A simpler example can be found with the oxygen molecule, O2:
O : : O or O = O
You will notice that each shared electron is counted twice—once for the left-hand oxygen
atom, and a second time for the right-hand oxygen atom. In this way, each oxygen atom
is surrounded by 8 electrons. A triple covalent bond, or triple bond, is a covalent bond
produced by the sharing of three pairs of electrons between two atoms. An excellent
simple example of this a nitrogen molecule, N2:
:N
N: or :N≡N:
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Unit I Notes: Chemical Bonding and Molecular Structure
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Double and triple bonds are often called multiple bonds, or multiple covalent bonds.
Double bonds will generally have higher bond energies (this means they’re stronger than
single bonds) and are shorter in length than single bonds. Triple bonds have even stronger
(higher bond energies yet) and are even shorter than double bond lengths. It is always
possible that nitrogen, oxygen, and carbon-containing molecules may have multiple
bonds between atoms in order to satisfy the octet rule. Hydrogen, however, will
always form single bonds, because it needs only one electron in order to complete its
valence shell.
We briefly discuss sigma () and pi () bonds in this unit as well. Examples of sigma
bonds include the overlap of 2 s orbitals (H2), the overlap of an s orbital and a p orbital
(HCl), the overlap of 2 p orbitals (Cl2), and the overlap of a p orbital with a hybrid orbital
(BeF2). The overlap occurs on a line that would connect the two nuclei involved in the
bond (along the bond axis).
The overlap for pi bonds occurs above and below (but not along) the line that would
connect the two nuclei involved in the bond and will always form from the overlap of p
orbitals. A double bond will always consist of 1  bond and 1  bond (see Figure 8.20 on
p. 235). A triple bond will always consist of 1  bond and 2  bonds (see Figure 8.21 on
p. 236). Both double and triple bonds are more rigid than single bonds, so as these
additional bonds are added, there is less and less rotational/vibrational motion along the
bond axis, and the bonds get shorter.  bonds are less strong than  bonds, so double- and
triple-bonded carbons tend to be more reactive than single-bonded atoms because the
electrons in the  bonds are farther from the nuclei of the two atoms, and therefore are
more weakly held. We call compounds with double and triple carbon atoms unsaturated
compounds.
We have already talked about the existence of polyatomic ions in our discussion of
compound naming and formula writing. These ions are made up of atoms of two or more
elements that are covalently bound. You can see the electron dot structure for the sulfite
ion on page 224. You should note that each atom manages to surround itself with 8
electrons to satisfy the octet rule, with the exception of hydrogen, which only requires 2
electrons. These polyatomic species will give up electrons to, or take electrons from an
outside source in order to achieve a positive or negative charge. On your ion charge
sheet, ammonium ion, NH 4 , is the only positively-charged polyatomic ion (see its
electron dot diagram in Figure 8.9 on page 223). All the rest have negative charges.
You should make sure you understand the dot structures that can be drawn to illustrate
the structure of polyatomic ions.
I-3 Reading
Scientists use two different theories to describe molecule geometry. The first theory is
called VSEPR (pronounced “Vesper”) theory. This is an acronym for the very long term
VALENCE-SHELL ELECTRON-PAIR REPULSION. We all know that like charges
repel, so it comes as no shock that electrons in the same molecule repel each other.
VSEPR theory states that repulsion between the sets of valence electrons surrounding the
central atom in a molecule cause these sets to be oriented as far from each other as
possible. We can use this theory to explain a whole host of molecule shapes, but we will
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Unit I Notes: Chemical Bonding and Molecular Structure
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focus on linear, bent, pyramidal, tetrahedral, trigonal planar, trigonal bipyramidal, and
octahedral.
Sometimes, there exists more than one possible Lewis structure. These equivalent
structures are called resonance structures, and they occur in many molecules (like O3
and benzene, C6H6) and polyatomic ions (like nitrate ion, NO3 , and carbonate ion, CO32  ).
The proper way to represent that molecule is to draw all equivalent structures, as can be
done with ozone and nitrate ion below:
(a) ozone resonance structures (18 e-)
(b) nitrate ion resonance structures (24 e-)
As mentioned before, there are molecules that violate the octet rule and they fall into
three different categories:

Incomplete octets (fewer than 8 electrons in the valence shell)
Some atoms can have a complete outer shell with less than 8 electrons—e.g.,
hydrogen is stable with two electrons, beryllium can be stable with only 4 valence
electrons (BeH2, etc.), and boron can be stable with only six valence electrons
(BF3, etc.)

Expanded octets (more than 8 electrons in the valence shell)
In molecules that have available d subshells (they are filling the third or higher
energy level), the central atom can have more than 8 valence electrons. Examples
include PCl5, SF4, and SF6. Take careful notes from our class discussion for
appropriate ways to represent these molecules with Lewis structures.

Odd numbers of electrons
Molecules almost always have an even number of electrons, which allow
electrons to be paired. There are some exceptions to this, most of them involving
nitrogen—NO, NO2, ClO2.
The simplest cases for us to discuss are ones in which the electrons in the valence shell
are all shared pairs. The easiest ones of these to understand is the 2-atom molecule
hydrogen gas, H2, or the 3-atom molecule BeH2. For H2, the shared pair is between the
two nuclei, and the repulsion within this pair is counterbalanced by the attraction to the
nuclei on either side of them. Because there is no central atom, this molecule will have a
linear shape because the only electrons involved will be between the nuclei:
H
H
BeH2 is a different case, because the central atom, Be, is capable of only forming two
covalent bonds. According to VSEPR theory, these two shared pairs will be oriented on
opposite sides of the beryllium nucleus in order to minimize repulsions:
H
Be H
It therefore makes sense that the shape of this molecule will also be linear, because the
angle formed between the bonds are 180.
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Unit I Notes: Chemical Bonding and Molecular Structure
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You may recall that we spoke earlier in the chapter about how boron doesn’t satisfy the
octet rule when it bonds covalently. It will surround itself with 3 shared pairs. When
boron bonds with fluorine, there will be 3 shared pairs that are oriented as far as possible
from each other. This leads to a kind of triangular structure that is all in the same plane
called trigonal planar. The angle between the bonds is 120:
When we move onto 4 shared pairs of electrons from 4 single bonds (like CH4), VSEPR
theory will no longer predict that all of the shared pairs will stay in the same plane. The
shared pairs actually form a 3-D geometry in which all 4 bonds (shared pairs) are
separated by angles of 109.5. This gives the molecule a tetrahedral shape. This
geometry is shown in Figure 8.16 (a) on page 232.
So far, things have been pretty straightforward (I hope), but they do get a bit more
complicated with the addition of unshared pairs of electrons, or regions of
concentrated negative charge, around the central atom. The first example molecule is
ammonia, NH3. Nitrogen has 5 electrons, and it needs 3 electrons in order to satisfy the
octet rule. % error calculations show that it forms 3 polar covalent bonds with each of 3
hydrogen atoms:
You will notice that after bonding there are 3 shared pairs of electrons(regions of
concentrated negative charge) and 1 unshared pair of electrons (region of concentrated
negtative charge) surrounding the central nitrogen atom. The unshared electron pair or
region does occupy space, so that one would assume that a tetrahedral shape would occur.
The text states, however, that the shape of the molecule is determined by only the
positions of the atoms, so a pyramid with a triangular base occupied by the hydrogen
atoms (the nitrogen atom is at the top of the pyramid) should be visualized. In addition,
unshared electron pairs repel more strongly than shared pairs of electrons, so the angle
between all of the bonds is now 107.5 because the unshared pair repels more strongly.
This situation, with 3 shared pairs and one unshared pair defines the pyramidal
geometry, as shown in Figure 8.16 (b) on page 232.
The last tetrahedral geometry that you need to know about is formed with 2 shared pairs
(regions of concentrated negtative charge), and 2 unshared pairs of electrons (regions of
concentrated negtative charge) around the central atom. This often occurs with molecules
having central atoms from Group 16. Water is the easiest example for this:
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Unit I Notes: Chemical Bonding and Molecular Structure
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The additional unshared electron pair (region of negative charge) also occupies space, so
that one would assume that a tetrahedral shape would occur. The text states again,
however, that the shape of the molecule is determined by only the positions of the atoms,
so a bent, or angular shape should be visualized. Once again, unshared electron pairs
repel more strongly than shared pairs of electron, so the angle between all of the bonds is
now further decreased to 104.5 because the 2 unshared pairs repel even more strongly.
This situation, with 2 shared pairs (regions) and 2 unshared pairs (regions) defines the
bent geometry. You can refer to Figure 8.17 on page 233.
Finally, in VSEPR theory, double and triple bonds are treated in the same way as single
bonds. This means that even though a double bond has 2 shared pairs, we treat it as
having only 1 region of concentrated negative charge when it comes to considering
molecular geometry. An example of this occurs with carbon and oxygen in CO2:
O=C=O
The central carbon atom has four shared pairs, but they are counted, according to VSEPR
theory as being only 2 shared regions of concentrated negative charge with no unshared
regions. This is identical to the BeH2 example from above, so this molecule will also be
linear.
Below is a handy table that I have taken from the Princeton Review AP chemistry text
(pp. 40-43) for telling the types of geometries that will form:
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Unit I Notes: Chemical Bonding and Molecular Structure
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If the central atom has 2 regions of concentrated negative charge then it has sp hybridization and its basic shape is linear.
Number of lone pairs
Molecular Geometry
0
linear
Examples
BeCl2, CO2
If the central atom has 3 regions of concentrated negative charge then it has sp2 hybridization and its basic
shape is trigonal planar.
Number of lone pairs
Molecular Geometry
0
trigonal planar
1
bent
Examples
BF3 , SO3 , NO3 , CO32
SO2
If the central atom has 4 regions of concentrated negative charge then it has sp3 hybridization and its basic
shape is tetrahedral.
Number of lone pairs
Molecular Geometry
0
tetrahedral
1
trigonal pyramidal
2
bent
Examples
CH 4 , NH 4 , ClO4 , SO42 , PO42
NH 3 , PCl3 , AsH 3 , SO32
H 2O, OF2 , NH 2
If the central atom has 5 regions of concentrated negative charge then it has sp3d hybridization and its basic
shape is trigonal bipyramidal.
Number of lone pairs
Molecular Geometry
Examples
0
trigonal bipyramidal
PCl5 , PF5
1
see-saw (irregular
tetrahedron)
SF4 , IF4
2
T-shaped
ClF3 , ICl3
3
linear
XeF2 , I 3
If the central atom has 6 regions of concentrated negative charge then it has sp3d2 hybridization and its basic
shape is octahedral.
Number of lone pairs
Molecular Geometry
0
octahedral
1
square pyramidal
2
square planar
Examples
SF6
BrF5 , IF5
XeF4 , ICl4
Remember to refer to the electron pair and molecular geometry pictures in our
text (Figures 8.12 on p. 229, and Figures 8.16, 8.17, and 8.18 on pages 232233) so that you have a clear understanding of how these molecular geometries
occur. You are only responsible for the situations that define the FIVE
GEOMETRIES (linear, tetrahedral, trigonal pyramidal, bent, and
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Unit I Notes: Chemical Bonding and Molecular Structure
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trigonal planar) FOR THE SOL, but are responsible for knowing the
trigonal bipyramidal and octahedral geometries for this unit and your
mid-term examination.
Now we come to the second theory that helps us to explain molecular geometry. You
might have already realized that atoms like carbon have an element dot structure that
looks like this:
C:
But we say that it somehow manages to look like this when it is forming covalent bonds:
C
What has happened to the s-orbital being filled before the p-orbitals? Scientists have
come up with the theory of hybridization to explain this. As you know, the s-orbital and
the p-orbital are at the same principal energy level, with the p-orbital being a bit higher.
We must visualize the mixing of two or more atomic orbital types of similar energies
on the same atom to produce new orbital types that have equal energies. You will
notice that carbon has a completely occupied 2s orbital and 2 electrons at the 2p sublevel:
_____ _____ _____ _____ _____
1s
2s
2p
In order to form the 4 identical single bonds that carbon forms with hydrogen when it
makes methane (CH4), the 2s orbital hybridizes with the three 2p orbitals to form 4
identical orbitals having a different shape called sp3 orbitals. Notice that the superscripts
(1 + 3) add up to the original number of orbitals:
_____ _____ _____ _____ _____
1s
2sp3
sp3 hybridization also explains the tetrahedral shape of the ammonia and water, as seen in
Figure 8.16 on page 232. A similar type of hybridization must also be used to explain
why beryllium is able to form two equal bonds when it only starts with a 1s22s2
configuration. The 2s electrons are hybridized to form two sp hybrid orbitals capable of
forming two single bonds. Any molecule that mimics that situation (two regions of
concentrated negative charge around the central atom, such as CO2) is also considered to
exhibit sp hybridization. The trigonal planar geometry discussed above for BF3 is seen to
exhibit sp2 hybridization. Any 4-atom geometry that mimics the 120 bond angle is
considered to exhibit the same hybridization. An example of this occurs, for example,
with the molecule, ethylene (sometimes called ethene) C2H4.
You will notice that double bonded carbon atoms exhibit sp2 hybridization with 120
bond angles, and triple bonded carbon atoms exhibit sp hybridization, which makes them
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linear molecules (180 bond angle). These types of hybridization, and their relationship
to the different geometries are summarized in the Princeton Review chart that I provided
above.
Delocalized  bonding occurs in molecules with more than one resonance structure.
This means that the  bonds can be shared throughout the molecule, as in the ring
compound benzene. When a carbon chain has alternating double and single bonds, the
molecule is said to contain a conjugated system. You can read pages 710-711 for
understanding of linear and ring compounds with these systems. For AP Chemistry you
will need to understand the representation of a resonance structure, and what it
means in terms of orbital overlap.
When there is more than one possible structure for the same chemical formula, we call
these different possibilities isomers. You have already probably seen the different types
of isomers in biology—they are called structural (linear vs. 3-D structures), positional
(a double or triple bond [or any atom other than carbon or hydrogen] is moved around
within the molecule), functional (when an atom or atoms other than carbon and hydrogen
are bonded in the molecule in 2 or more ways), and geometric (different arrangements of
atoms around a double bond) isomers. Examples of these are shown and discussed on
pages 704-706.
I-4 Reading
The last part of our discussion focuses on the fact that some covalent compounds are not
just molecules hanging out in the same area—they form what is called covalent network
bonding, in which many identical molecules are bound together by forces acting between
the molecules. These are intermolecular forces, and they are discussed at the end of the
unit. There are also many covalently bonded compounds that don’t contain individual
molecules, but instead are continuous 3-Dimensional networks of covalently bonded
atoms. Polymers (plastics) are a good example of this.
Properties of molecular compounds are determined by whether the individual molecules
are polar (uneven distribution of charge within a molecule) or nonpolar (even distribution
of charge within a molecule). This molecule polarity is influence by two factors: the
polarity of individual bonds within a molecule, and the overall geometry of the molecule.
We will now discuss in how molecule polarity affects overall intermolecular forces.
Strangely enough, it is possible to have a molecule with individual polar covalent bonds,
but for the molecule as a whole to be nonpolar. Here’s how it arises. Let’s say that we
have a molecule that forms from carbon and hydrogen. %IC tells us that an individual
bond formed between a carbon atom and a hydrogen atom will be slightly positive on the
hydrogen side of the bond, and slightly negative on the carbon side of the bond. A small
separation of charge, or dipole, has formed. Dipoles occur whenever equal and opposite
charges are separated by a distance. The size of the dipole is measured by the dipole
moment. The larger the dipole moment, the more polar the molecule will be. You can see
by the mathematical definition of the dipole moment:
=Qd
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that the greater the charge, Q, the greater will be the dipole moment. Similarly, the
further distance, d, apart the opposite charges are, the greater will be the dipole moment
as well (this is a directly proportional relationship). We indicate polarity in a molecule
with a  - by the atom with the higher electronegativity, and a  + next to the atom with
the lower electronegativity, as shown in the HF molecule:


H F
We also know that carbon, after hybridization, is capable of forming 4 single bonds with
4 Hydrogen atoms. As we have already discussed, the resulting molecule shape is
tetrahedral. As long as all of the atoms bonded to the central carbon atom are the same,
the polarity of each of the bonds formed will be identical. In addition to that, this
molecule exhibits radial symmetry, which means that the distribution of the electrons
around the central atom is uniform—the same in every direction. This means that all of
the dipoles effectively cancel each other out so that there is no net polarity on the
final molecule.
Pyramidal and bent shapes do not exhibit this radial symmetry, so if their individual
bonds are polar covalent, the final molecules will also be polar. Linear 2-atom molecules
will be nonpolar molecules only if the bond between the atoms is nonpolar. Linear 3atom molecules can exhibit radial symmetry if the central atom is bonded on both sides
with the same element (e.g. BeH2)
We will have lots of practice with molecule geometry and molecule polarity with the
VSEPR lab activity.
We just finished talking about polar and nonpolar molecules. How will these affect
intermolecular attractions? First, let me state that any two objects that have mass will
experience a “gravitational” attraction to each other. Some experience stronger attractions
than others. We do know that solid molecules are moving much slower than liquid
molecules, which are moving slower than gas molecules of the same substance. As a
substance is heated the energy being added is converted to kinetic energy so that the
energy is sufficient to overcome the attractive forces. This is why substances melt and
then vaporize as more energy is added to the system. We also know that different
substances will have different melting and boiling points. The only possible explanation
for this is that the molecules of each substance must experience different levels of
attractive forces between them. The stronger the attraction, the more energy will be
needed to overcome it.
There are different types of intermolecular forces of varying strength, but in general they
are weaker than ionic or covalent bonds. In general BP’s of metals and ionic compounds
are higher than those for molecular compounds. The more polar the molecule is,
however, the greater will be the intermolecular forces that we discuss below.
Van der Waals Forces are weak intermolecular forces that can be broken down into two
classifications:
Polar molecules are dipoles—that is, they are electrically neutral as a whole, but they
exhibit a separation of charge over a small distance. When two polar molecules are near
Intensified Chemistry
Unit I Notes: Chemical Bonding and Molecular Structure
Page 13 of 13
each other, the positive end of one will be attracted to the negative end of the other
through dipole-dipole forces, as shown in Figure 8.25 on page 240. The presence of
these forces cause MP’s and BP’s to be higher than if they weren’t present.
A polar molecule can also cause a nonpolar molecule to temporarily form a dipole by
attracting or repelling its electrons. These are called dipole-induced dipole forces, and
they are weaker intermolecular forces than ordinary dipole-dipole forces.
The second type of weak intermolecular forces that we need to consider is London
Dispersion Forces. Because electrons are in constant motion, it is possible that any
moment, a spontaneous dipole could arise from an asymmetrical distribution of
electrons in a normally nonpolar molecule like a noble gas or H2. This temporary
dipole can induce a dipole in another molecule, and there will be attraction between them.
The more electrons there are in a molecule, the greater the chance that London
Dispersion Forces will be important. This means that the greater the molar mass of the
substance is, the greater will be the London Dispersion Force contribution to the
properties of a substance.
A very strong type of dipole-dipole force is hydrogen bonding. It occurs in substances
where hydrogen is bonded to extremely electronegative elements (O, N, and F). The
molecules that form are highly polar, and the strongly positive charge hydrogen atom is
so small that it can come very close to an unshared electron pair on a nearby molecule.
The resulting attraction leads to a very strong intermolecular attraction. We represent
hydrogen bonds with dotted lines.
We can compare the properties of ionic and molecular compounds. Both ionic and
molecular bonding forces are strong, but the forces between molecules (not necessarily
within them) are much weaker than the forces of ionic bonding. This difference in
forces between the formula units of ionic compounds and the forces between covalent
molecules actually cause big differences in the bulk (read this is as large sample size)
properties of the two types of compounds.
Melting point (MP), boiling point (BP), and hardness of compounds all depend on how
strongly the basic units are attracted to each other. The greater the attraction, the higher
these values will be. Because the forces between molecules are not very strong,
molecular compounds will melt at low temperature, and ionic compounds will melt and
boil at very high temperatures compared to molecular compounds. Molecular compounds
will also evaporate more readily than ionic compounds at room temperature because of
lower intermolecular attractions.
Ionic compounds are hard (think of ceramics), but they are brittle. This is because even a
slight shift in the location of ions can cause overwhelming repulsive forces that make the
solid come apart easily. Ionic compounds are only good conductors when they are in the
molten (liquid) state, or when they are dissolved in water. This is the only time when the
ions are free to move and carry electrical current. When the ions dissolve in water they
are separated and surrounded by water molecules in a process called solvation. We will
talk more about this later on when we discuss solutions and solubility.
Table 8.4 on page 244 summarizes ionic and covalent compound properties due to intraand intermolecular forces of attraction.