Topic 6: Chemical Bonding &
Molecular Geometry
Chemical Bonding
(Chapter 6 in Modern Chemistry)
Atoms seldom exist as independent particles in nature. Most substances
consist of combinations of atoms that are held together by chemical bonds. A chemical bond is
a mutual electrical attraction between the nuclei and valence electrons of different atoms that
binds the atoms together. Why do atoms make bonds? It turns out that most atoms are less
stable existing by themselves than when they are combined. By bonding with each other, atoms
decrease in potential energy, thereby creating more stable arrangements of matter.
When atoms bond, their valence electrons are redistributed in way that make the atoms more
stable. The way in which the electrons are redistributed determines the type of bonding. In
Topic 4 you learned the main-group metals tend to lose electrons to form positive ions, or
cations. Nonmetals tend to gain electrons to form negative ions, or anions. Chemical bonding
that results from the electrical attraction between cations and anions is called ionic bonding. If
a bond is purely ionic, an atom will completely give up electron(s) to another atom.
In contrast, atoms joined by covalent bonding share electrons. Covalent bonding results from
the sharing of electron pairs between two atoms. In a purely covalent bond, the shared electrons
are “owned” equally by the two bonded atoms.
an example of an ionic bond
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an example of a covalent bond
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Ionic or Covalent?
Bonding between atoms is rarely purely ionic or purely covalent. It usually falls somewhere
between these two extremes, depending on how strongly the atoms of each element attract
electrons. Recall that electronegativity is a measure of an atom’s ability to attract electrons. To
determine whether a bond is ionic or covalent you have to calculate the difference in
electronegativities of the two atoms involved.
An electronegativity difference greater than 1.67 is referred to as an ionic bond.
Electronegativity differences of 1.67 or less have an ionic
bond character of 50% or less. These compounds are
typically classified as covalent. Bonding between two atoms
of the same element is completely covalent. This is called a
nonpolar-covalent bond. This is a covalent bond in which
the bonding electrons are shared equally by the bonded
atoms, resulting in a balanced distribution of electrical
charge. Bonds having only 0% to 5% ionic character, or an
electronegativity difference equal to or less than 0.3, are
considered nonpolar-covalent.
Bonds having an ionic character between 5% and 50%, or
with corresponding electronegativity differences of 0.3 to
1.67, are classified as polar-covalent. A polar-covalent
bond is a covalent bond in which the bonded atoms have an
unequal attraction for the shared electrons.
3.3
100%
Ionic
1.67
50%
Polar -covalent
0.3
Nonpolar-covalent
0
5%
0%
Here are some example of determining bond character based on the electronegativities from the
periodic table on the next page.
Bonding Pair
Electronegativity
difference
Bond type
Li and F
3.98 - 0.98 = 3.00
ionic
Cu and S
2.58 – 1.90 = 0.68
polar-covalent
I and Br
2.96 -2.66 = 0.30
nonpolar-covalent
In summary, subtract the electronegativities, if:
Greater than 1.67
Above 0.3 to 1.67
0.3 or less
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ionic
polar-covalent
nonpolar-covalent
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Electronegativities of the elements
→ Atomic radius decreases → Ionization energy increases → Electronegativity increases →
Group
1
2
3
4
5 6
7
8
9 10 11 12 13 14 15 16 17 18
(vertical)
Period
(horizontal)
H
He
1
2.20
Li Be
B C N O F Ne
2
0.98 1.57
2.04 2.55 3.04 3.44 3.98
Na Mg
Al Si P S Cl Ar
3
0.93 1.31
1.61 1.90 2.19 2.58 3.16
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
4
0.82 1.00 1.36 1.54 1.63 1.66 1.55 1.83 1.88 1.91 1.90 1.65 1.81 2.01 2.18 2.55 2.96 3.00
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe
5
0.82 0.95 1.22 1.33 1.6 2.16 1.9 2.2 2.28 2.20 1.93 1.69 1.78 1.96 2.05 2.1 2.66 2.60
Cs Ba * Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
6
0.79 0.89
1.3 1.5 2.36 1.9 2.2 2.20 2.28 2.54 2.00 1.62 2.33 2.02 2.0 2.2 2.2
Fr Ra ** Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uus Uuo
7
0.7 0.9
Lanthanoids
Actinoids
*
La
1.1
** Ac
1.1
Ce Pr Nd Pm Sm Eu Gd
1.12 1.13 1.14 1.13 1.17 1.2 1.2
Th Pa U Np Pu Am Cm
1.3 1.5 1.38 1.36 1.28 1.13 1.28
Tb
1.1
Bk
1.3
Dy Ho Er Tm
1.22 1.23 1.24 1.25
Cf Es Fm Md
1.3 1.3 1.3 1.3
Yb
1.1
No
1.3
Lu
1.27
Lr
1.3
Periodic table of electronegativity using the Pauling scale
Please note, if you do not have the electronegativities, ionic compounds are generally made up of
elements that are far apart on the periodic table. For example: K & Cl make an ionic bond. Sr
& Br make an ionic bond. S & O make a covalent bond. (Watch out for hydrogen; even though
it is in group 1, it has a high electronegativity.)
Task 6a
1. Using the electronegativity chart, determine the type of bond between the atoms in the
following compounds.
b)
c)
d)
a. AgCl a)
K - 0.82
Br
2.96
H - 2.2
K - 0.82
b. K2O Ag - 1.93
Br - 2.96
Cl - 3.16
Cl - 3.16
O - 3.44
2.96 - 2.96 = 0
3.16 - 2.2 = 0.96
c. Br2
3.16 - 1.93 = 1.23
3.44 - .82 - .82 = 1.8
Bond
—>
Nonpolar-covalent
Bond —> Polar-covalent
d. HCl Bond —> Polar-covalent Bond —> Polar-covalent
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2. Only using the periodic table, determine if the bonds below are covalent or ionic. Do not
use the electronegativity chart.
a. XeCl6 Covalent
b. CsF Ionic
c. MgCl2 Ionic
d. NO2 Covalent
Ionic Bonding and Ionic Compounds
An ionic compound is composed of positive and negative ions that are combined so that the
numbers of positive and negative charges are equal. The chemical formula of an ionic
compound shows the ratio of ions present in a sample of any size. A formula unit is the
simplest collection of atoms from which an ionic compound’s formula can be established. For
example, one formula unit of sodium chloride, NaCl, is one sodium cation plus one chloride
anion. (In the naming of a monatomic anion, the ending of the element’s name is replaced with –ide.) The ratio of ions in a formula unit depends on the charges of the ions combined. They
must achieve electrical neutrality. In other words, the sum of the charges must equal zero.
For example: calcium fluoride
calcium has a 2+ charge, Ca2+
fluorine has a 1- charge, FIn order for their charges to equal zero, there has to be two fluoride ions for every
calcium ion.
CaF2
The Formation of Ionic Compounds
Consider that a sodium atom and a chlorine atom are approaching each other. The two atoms are
neutral and have one and seven valence electrons, respectively.
Sodium loses its electron to chlorine forming Na + and Cl-.
Sodium atom
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+
Chlorine atom
Sodium ion
+
Chloride ion
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Covalent Bonding
Most substances are composed of molecules. A molecule is a neutral group of atoms that are
held together by covalent bonds. A molecule may consist of two of more atoms of the same
element, as in oxygen, or of two or more different atoms, as in water or sugar.
Oxygen molecule,
Water molecule,
O2
Sucrose molecule,
H2O
C12H22O11
A chemical compound whose simplest units are molecules is called a molecular compound.
The composition of a compound is given by its chemical formula. A chemical formula
indicates the relative numbers of atoms of each kind in a chemical compound by using atomic
symbols and numerical subscripts. The chemical formula of a molecular compound is referred
to as a molecular formula. A molecular formula shows the types and numbers of atoms
combined in a single molecule of a molecular compound. The molecular formula for water, for
example, is H2O, which reflects the fact that a single water molecule consists of one oxygen
atom joined by separate covalent bonds to two hydrogen atoms. A molecule of oxygen, O 2, is an
example of a diatomic molecule. A diatomic molecule is a molecule containing only two atoms.
Formation of a Covalent Bond
Remember that nature favors chemical bonding because most atoms have lower potential energy
when they are bonded to other atoms than they have when they are not bonded.
Using the example of two hydrogen atoms, if they are separated by enough distance, they will
not influence each other. However, if the two hydrogen atoms approach each other, there will
come to a position in which the nucleus of one hydrogen atoms attracts the electron from the
other hydrogen atom. The attraction corresponds to a decrease in the total potential energy of the
atoms. At the same time the electrons of the two hydrogen atoms are repelling each other. The
protons in the nuclei of the two hydrogen atoms are also repelling each other. The repulsion
results in an increase in potential energy. The relative strength of attraction and repulsion
between the charged particles is dependent on the distance separating the two hydrogen atoms.
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Here the arrows indicate the attractive and repulsive forces
between the electrons and nuclei of two hydrogen atoms.
Attraction (red) between particles corresponds to a decrease in
potential energy of the atoms, while repulsion (blue)
corresponds to an increase.
The attractive force dominates and continues to pull the two hydrogen atoms closer together until
they get to a distance at which the repulsion between like charges equals the attraction between
opposite charges. At this position the two hydrogen atoms (now a hydrogen molecule, H 2) have
their minimum potential energy and are close enough to share electrons. They are now
covalently
bonded.
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Hydrogen
atoms
Hydrogen
molecule
All individual hydrogen atoms contain a single, unpaired electron in a 1s atomic orbital. When
two hydrogen atoms form a molecule, they share electrons in a covalent bond. The sharing of
the electrons allows each atom to have the stable electron configuration of helium, 1s2. The
tendency is for atoms to achieve a noble-gas configuration by sharing electrons.
The Octet Rule
Unlike other atoms, the noble-gas atoms exist independently in nature. They possess a minimum
of energy existing on their own because of the special stability of their electron configuration.
This stability results from the fact that, with the exception of helium and its two electrons in a
completely filled outer shell, the noble-gas atoms’ outer s and p orbitals are completely filled by
a total of eight electrons.
Other main-group atoms can effectively fill their outermost s and p orbitals with electrons by
sharing electrons through covalent bonding. Such bond formations follow the octet rule.
Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons,
has an octet of electrons init highest occupied energy level. Here is an example of the covalent
bonding of hydrogen and fluorine to make hydrofluoric acid, HF.
H
H
F
F
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Exceptions to the Octet Rule
Most main-group elements tend to form covalent bonds according to the octet rule, However,
there are exceptions. Hydrogen forms bonds in which there are only two electrons. Boron, B,
and aluminum, Al, has just three valence electrons so it makes bonds surrounded by 6 electrons.
Other atoms can be surrounded by more than eight electrons. Examples of elements that make
bonds with expanded valences are phosphorus, P, in PF5 and sulfur, S, in SF6.
Task 6b
1. Use orbital notation to illustrate the bonding in the chlorine molecule, Cl 2.
2. Describe the general location of the electrons in a covalent bond.
Metallic Bonding
Chemical bonding is different in metals than it is in ionic or molecular compounds. The highest
energy levels of most metal atoms are occupied by very few electrons. The properties of metals
are due to the highly mobile valence electrons of the atoms that make up a metal. The highest
energy levels of most metal atoms are occupied by very few electrons. They have many vacant
orbitals. The vacant orbitals in the atoms’ outer energy levels overlap. This overlapping of orbitals allows the outer electrons of the atoms to roam freely throughout the entire metal. The
electrons are delocalized, which means that they do not belong to any one atom but move freely
about the metal’s network of empty atomic orbitals. These mobile electrons form a sea of
electrons around the metal atoms, which are packed together in a crystal lattice. The chemical
bonding that results from the attraction between metal atoms and the surrounding sea of
electrons is called metallic bonding.
Here are two diagrams to help you understand how the electrons surround the cation in a metallic
bond
.
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Van der Waals force
Van der Waals forces are a group of weak intermolecular forces that vary in strength but are
generally weaker than bonds (ionic, covalent, and metallic). We will discuss hydrogen bonding,
dipole-dipole interactions, and London dispersion forces (LDF).
Hydrogen Bonding
Some hydrogen-containing compounds, such as hydrogen fluoride, water, and ammonia, have
unusually high boiling points. This is explained by the presence of a particularly strong type of
dipole-dipole force. In compounds that contain hydrogen and a very electronegative element (N,
O, F) an intermolecular attraction occurs between the molecules. Note that this is not a bond
between atoms within the molecule, but an attraction among the molecules. Below, the dotted
lines represent the hydrogen bonds between water molecules. This allows water to have
relatively high melting and boiling points for hydrogen compounds that bond with group 16
nonmetals. This also explains why ice expands and floats.
Dipole-Dipole Attractions
A dipole is a molecule or a part of a molecule that contains both positively and negatively
charged regions. For example, the molecule, HCl is made with hydrogen
and the very electronegative atom chlorine. The electrons in this molecule
tend to gather around the chlorine. This makes the hydrogen end more
positively charged ( read as partially positive) and the chlorine end more
negative ( -, read as partially negative). Each end is a dipole.
H
Cl
+
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-
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That means that when two hydrogen chloride molecules get
close together and they are oriented correctly, the different
dipoles will be attracted to each other.
London Dispersion Forces
There is some type of intermolecular force among all atoms and molecules, even noble gases.
These are the weakest of the van der Waals forces and are called the London dispersion forces.
London dispersion forces (LDF) are intermolecular attractions resulting from the constant
motion of electrons and the creation of instantaneous dipoles. The more electrons there are the
greater the London dispersion forces. So it goes to reason that the bigger the atomic mass or
molecular mass the greater the London dispersion forces. That is why boiling points generally
increase down a group on the periodic table.
Summary of Bonding and their relationship to properties
Ionic
Covalent
Metallic
(Giant lattice)
(Individual molecules)
(Mixtures of Metals)
Small groups of atoms
covalently bonded together
by sharing electrons
Close packed array of
atoms (ions) with "sea" of
free moving electrons
Metal + Non-Metal
ex: KCl, MgF2, Na2SO4, etc.
Non-Metal + Non-Metal
ex: CO2, PCl5, etc.
Metals ex: Na, Al, Au,
Stainless Steel (Fe/C/Cr),
Bronze (Cu/Sn), Brass (Cu/Zn)
Strong Bonds
Strong bonds within
molecules, but weak
between molecules
Strong (but
flexible) bonds
low m.p/b.p., often liquids
or gases at RT;
LDF- (induced dipoles)
Dipole - perm. dipoles)
H-Bonds - (H - N/O/F)
poor conductors
high m.p/b.p. good
conductors of electricity ("sea'
of electrons) and heat (close
packed), malleable (can be
shaped), ductile (drawn into
thin wires), luster (shiny)
+ve ion and -ve ion formed
via transfer of electrons held
together in giant lattice with
strong electrostatic
interaction
high m.p./b.p.
poor conductors of
electricity when solid (ions
not free to move), good
when liquid or in solution
(dissolved)
will dissolve in polar solvents
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Task 6c
1. Compound B has lower melting and boiling points than compound A. At the same
temperature, compound B vaporizes faster than compound A. If one of these compounds
is ionic and the other is molecular, which would you expect to be molecular? Ionic?
Explain your reasoning.
2. Analyzing Data. The melting points for the compounds Li 2S, Rb2S, and K2S are 900oC,
530oC, and 840oC, respectively. List these three compounds in order of increasing lattice
energy.
3. Explain why most metals are malleable and ductile but ionic crystals are not.
4. Explain why metals are good electrical conductors.
5. What is the difference between a formula unit and a molecule.
Lewis Structures
In Topic 3, you learned about electron dot diagrams for elements. These can be used to represent
molecules also. When representing molecules, they are called Lewis structures. Lewis
structures are formulas in which atomic symbols represent nuclei and inner-shell electrons, dotpairs or dashes between two atomic symbols represent electron pairs in covalent bonds, and dots
adjacent to only one atomic symbol represent unshared electrons. If only the shared pairs
(bonds) are written using dashes, then this will be a structural formula. A structural formula
indicates the kind, number, arrangement, and bonds but not the unshared pairs of the atoms in a
molecule.
It is important to learn how to draw the Lewis structures of molecules and to predict the
molecular geometry of the molecule. To predict the geometry, you have to consider shared
(bonded) electrons and lone (nonbonded) electron pairs surrounding the central atom. To do this,
we will use the VSEPR theory.
VSEPR stands for “valence shell electron pair repulsion. That means that the electron pairs around the atoms (usually the central atom) repel each other to affect the shape of the molecule.
VSEPR theory states that the repulsion between the sets of valence-level electrons surrounding
an atom causes these sets to be oriented as far apart as possible. Unshared electrons repel the
most. Remember these molecules are 3-D even though we draw them in 2-D.
Drawing Lewis structures
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Here are the steps for writing a Lewis structure for compounds for which you do not know the
formula.
1. Calculate the total number of valence electrons taking into account any charges. (add
electrons for negative charges and subtract electrons for positive charges)
2. Decide which element is the central atom. Usually this is obvious; if in doubt it will be the
least electronegative atom (except for H: it only wants 2 electrons). Put that element in the
center and add the other elements around the central atom using lines to represent bonding pairs
of electrons.
3. Arrange the remaining electrons to complete the octet of the terminal atoms by placing pairs
of dots around the atoms. Place any remaining electrons (dots) on the central atom, if necessary
expanding the octet.
4. If the central atom lacks an octet, form multiple bonds (double or triple bonds) by converting
non-bonding electrons on terminal atoms into bonding pairs. (Sometimes atoms will remain
electron deficient).
For example:
CCl4
Valence electrons = 4 + 7(4) = 32
H2O
Valence electrons = 1(2) + 6 = 8
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Now practice with your teacher the following:
NH3
HF
PF5
NH4+
PCl6-
CO32-
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Task 6d
1. Draw a Lewis structure for each of the following molecules. You will come back to this
section for later tasks, so be sure to be neat and orderly.
a. SCl2
b. PI3
c. Cl2O
d. NH2Cl
N is central atom
T
e. SiCl3Br
Si is central atom
f. ONCl
O is central atom
g. SO42-
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h. ClO2-
i. BeCl2
You can also use VSEPR theory to predict the shapes of a molecule. Here is a link that might
help you understand the shapes.
http://library.thinkquest.org/C006669/data/Chem/bonding/shapes.html
You will need to learn the chart on the next page.
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Task 6e
1. Using the Lewis structures that you drew in 6d, determine the molecular geometry for
each.
Hybridization
VSEPR theory helps determine the shapes of a molecule but it does not reveal the relationship
between a molecule’s geometry and the orbitals occupied by its bonding electrons. To explain how the orbitals of an atom become rearranged when the atom forms covalent bonds, a different
model is used. This model is called hybridization, which is the mixing of two or more atomic
orbitals of similar energies on the same atom to produce new hybrid atomic orbitals of equal
energies.
Methane, CH4, provides a good example of how hybridization is used to explain the geometry of
molecular orbitals. The orbital notation for carbon’s valence electrons is 2s 2 2p2. We know form
experiments that a methane molecule has tetrahedral geometry. How does carbon form four
equivalent, tetrahedrally arranged covalent bonds? The one s orbital and the three p orbital
hybridize to form four new, identical orbitals called sp3 orbitals.
2p
sp3
2s
Carbon’s orbitals
Before hybridization
Carbon’s orbitals after
sp3 hybridization
Hybrid orbitals are orbitals of equal energy produced by the combination of two or more
orbitals on the same atom. The number of hybrid orbitals produced equals the number of
orbitals that have combined. A molecule with 3 bonding areas will be sp2, while a molecule with
4 bonding areas will be sp3. Trigonal bipyramidal molecules will be dsp3 and octahedral
molecules will be d2sp3.
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Polarity & Dipole forces
The strongest intermolecular forces exist between polar molecules. Polar molecules act as tiny
dipoles because of their uneven charge distribution. A dipole is created by equal but opposite
charges that are separated by a short distance. The direction of the dipole is from the dipole’s positive pole to its negative pole. A dipole is represented by an arrow with a head pointing
toward the negative pole and a crossed tail situated at the positive pole. The negative pole will
be the atom that is the most electronegative. For example, hydrochloric acid, HCl:
H__Cl
This is a polar bond and a polar molecule because the charges are unevenly distributed. If the
charge distribution is evenly distributed, the molecule will be nonpolar. For example CH4:
H
H __ C __ H
H
Notice that all the dipole point toward the C. Here the charge is evenly distributed. Therefore
the molecule is nonpolar even though the individual bonds are polar.
Task 6f
1. Go back to Task 6d and label the hybridization and the polarity of each molecule.
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