W5-1-VSEPR and Hybridization

Week No. 5-1 CHM61-101 Chemistry I Walailak University

VSEPR and Hybridization

Assoc. Prof. Dr. Phimphaka Harding
School of Science

Lesson objectives

• Define molecular geometry
• Define valence-shell electron-pair repulsion model
• Note the difference between the arrangement of electron pairs about a central atom and molecular geometry
• Define dipole moment
• Define σ (sigma) bond.
• Define π ( pi) bond

Learning outcomes Keywords

Students should be able to: Molecular geometry, Dipole moment, Sigma bond, Pi bond
• Use the Valence Shell Electron Pair Repulsion (VSEPR) model to predict

the geometries of molecules and polyatomic ions, including those with
more than one central atom
• Predict the geometries of simple polyatomic molecules using VSEPR
model
• Use electronegativity to predict molecular polarity
• Describe single and multi covalent bonds using valence bond theory
• Use the concept of hybridge orbitals to explain the geometries of simple
polyatomic molecules

VSEPR Theory

VSEPR Theory

In 1956, British scientists R.J. Gillespie and R.S. Nyholm recognized
that the current model for explaining bond angles did not work
well. The theory at that time relied on hybrid orbitals to explain all
aspects of bonding. The problem was that the theory gave
incorrect prediction of bond angles for many compounds. They
developed a new approach based on earlier work by other
scientists that incorporated a consideration of electron pairs in
predicting three-dimensional structure.

The valence shell is the outermost electron-occupied shell of an
atom. The valence shell holds the electrons that are involved in
bonding and are the electrons shown in a Lewis structure. The
acronym VSEPR stands for the valence-shell electron pair repulsion
model. The model states that electron pairs will repel each other
such that the shape of the molecule will adjust so that the valence
electron-pairs stay as far apart from each other as possible.
Molecules can be systematically classified according to the number
of bonding pairs of electrons as well as the number of nonbonding
or lone pairs around the central atom. For the purposes of the
VSEPR model, a double or triple bond is no different in terms of
repulsion than a single bond.

Summary

• VSEPR theory allows more accurate predictions of molecular shape.

Final wrap-up

Review
1. Who did the major work in developing the VSEPR theory?
2. Where are the electrons that are involved in bonding?
3. What is the basic idea behind the VSEPR theory?

PaMttoelrencsulaanrdShtraepnedss: iNnothLeonpeerPiaoidrsicotnabClentral Atom

How does an electroscope work? AB2: Beryllium hydride (BeH2)

An electroscope is a device used to study charge. When a positively Beryllium hydride consists of a central beryllium atom with two single bonds to
charged object (the rod) nears the upper post, electrons flow to the top hydrogen atoms. Recall that it violates the octet rule.
of the jar leaving the two gold leaves postivley charged. The leaves repel H-Be-H
each other since both hold postive, like charges. The VSEPR theory says
that electron pairs, also a set of like charges, will repel each other such According to the requirement that electron pairs maximize their distance from one
that the shape of the molecule will adjust so that the valence electron- another, the two bonding pairs in the BeH2 molecules will arrange themselves on
pairs stay as far apart from each other as possible. directly opposite sides of the central Be atom. The resulting geometry is a linear
molecule, shown in the Figure below in a “ball and stick” model.
Central Atom with No Lone Pairs
The bond angle from H-Be-H is 180° because of its linear geometry.
In order to easily understand the types of molecules possible, we will use
a simple system to identify the parts of any molecule. Carbon dioxide is another example of a molecule which falls under the AB2 category.
Its Lewis structure consists of double bonds between the central carbon and the
A = central atom in a molecule oxygen atoms (see Figure below).
B = atoms surrounding the central atom
Subscripts after the B will denote the number of B atoms that are The repulsion between the two groups of four electrons (two pairs) is no different than
bonded to the central A atom. For example, AB4 is a molecule with a the repulsion of the two groups of two electrons (one pair) in the BeH2 molecule.
central atom surrounded by four covalently bonded atoms. Again, it Carbon dioxide is also linear (see Figure below).
does not matter if those bonds are single, double, or triple bonds.

PaMttoelrencsulaanrdShtraepnedss: iNnothLeonpeerPiaoidrsicotnabClentral Atom

AB3: Boron Trifluoride (BF3) AB4: Methane (CH4)

Boron trifluoride consists of a central boron atom with three single bonds Methane is an organic compound that is the primary component of natural gas. Its
structure consists of a central carbon atom with four single bonds to hydrogen atoms
to fluorine atoms(see Figure below). The boron atom also has an (see Figure below). In order to maximize their distance from one another, the four
groups of bonding electrons do not lie in the same plane. Instead, each of the
incomplete octet. hydrogen atoms lies at the corners of a geometrical shape called a tetrahedron. The
carbon atom is at the center of the tetrahedron. Each face of a tetrahedron is an
equilateral triangle.

The geometry of the BF3 molecule is called trigonal planar (see Figure The molecular geometry of the methane molecule is tetrahedral (see Figure below).
below). The fluorine atoms are positioned at the vertices of an The H-C-H bond angles are 109.5°, which is larger than the 90° that they would be if the
molecule was planar. When drawing a structural formula for a molecule such as
equilateral triangle. The F-B-F angle is 120° and all four atoms lie in the methane, it is advantageous to be able to indicate the three-dimensional character of
its shape. The structural formula below is called a perspective drawing. The dotted
same plane. line bond is to be visualized as receding into the page, while the solid triangle bond is
to be visualized as coming out of the page.

Summary

• Electron pairs repel each other and influence bond angles and molecular shape.

Final wrap-up

Review
1. What are the bond angles in carbon dioxide?
2. What molecule has bond angles of 109.5°?
3. What is the geometry of the BF3 molecule?

PaMttoelrencsulaanrdShtraepnedss: iLnotnheePpaeirr(iso)doicntCaebnletral Atom

How can all these clothes fit into such a small space? AB3E: Ammonia, NH3

When we travel, we often take a lot more stuff than we need. Trying to The ammonia molecule contains three single bonds and one lone pair on the central
fit it all in a suitcase can be a real challenge. We may have to repack or nitrogen atom (see Figure below).
just squeeze it all in. Atoms often have to rearrange where the electrons
are in order to create a more stable structure. The domain geometry for a molecule with four electron pairs is tetrahedral, as was
seen with CH4. In the ammonia molecule, one of the electron pairs is a lone pair rather
Central Atom with One or More Lone Pairs than a bonding pair. The molecular geometry of NH3 is called trigonal pyramidal (see
Figure below).
The molecular geometries of molecules change when the central atom
has one or more lone pairs of electrons. The total number of electron Recall that the bond angle in the tetrahedral CH4 molecule is 109.5°. Again, the
pairs, both bonding pairs and lone pairs, leads to what is called the replacement of one of the bonded electron pairs with a lone pair compresses the angle
electron domain geometry. When one or more of the bonding pairs of slightly. The H-N-H angle is approximately 107°.
electrons is replaced with a lone pair, the molecular geometry (actual
shape) of the molecule is altered. In keeping with the A and B symbols
established in the previous section, we will use E to represent a lone pair
on the central atom (A). A subscript will be used when there is more
than one lone pair. Lone pairs on the surrounding atoms (B) do not affect
the geometry.

CaMtioolnecFuolramr Sahtiaopnes: Lone Pair(s) on Central Atom

AB2E2: Water, H2O The sulfur atom has five electron groups around it, which corresponds to the
trigonal bipyramidal domain geometry, as in PCl5 (see Figure below). Recall
A water molecule consists of two bonding pairs and two lone pairs (see Figure that the trigonal bipyramidal geometry has three equatorial atoms and two
below). axial atoms attached to the central atom. Because of the greater repulsion
of a lone pair, it is one of the equatorial atoms that are replaced by a lone
As for methane and ammonia, the domain geometry for a molecule with four pair. The geometry of the molecule is called a distorted tetrahedron or
electron pairs is tetrahedral. In the water molecule, two of the electron pairs seesaw.
are lone pairs rather than bonding pairs. The molecular geometry of the water
molecule is bent. The H-O-H bond angle is 104.5°, which is smaller than the
bond angle in NH3 (see Figure below).

AB4E: Sulfur Tetrafluoride, SF4

The Lewis structure for SF4 contains four single bonds and a lone pair on the
sulfur atom (see Figure below).

CaMtioolnecFuolramr Sahtiaopnes: Lone Pair(s) on Central Atom

Total Number of Number of Bonding Number of Lone Electron Domain Molecular Examples
Electron Pairs Pairs Pairs Geometry Geometry
3 2 1 O3
4 3 1 trigonal planar bent NH3
4 2 2 H2O
5 4 1 tetrahedral trigonal pyramidal SF4

5 3 2 tetrahedral Bent ClF3
5 2 3 trigonal bipyramidal I3-
6 5 1 distorted BrF5
6 4 2 trigonal bipyramidal tetrahedron XeF4

(seesaw)

T-shaped

trigonal bipyramidal Linear

Octahedral square pyramidal

Octahedral square planar

Summary

• The presence of lone pair electrons influences the three-dimensional shape of the molecule.

Final wrap-up

Review
1. Why does water have a bent geometry?
2. Why is ammonia not a planar molecule?
3. How would we write the configuration for xenon tetrafluoride using the ABE system?

PBaotntderPnoslanridtytrends in the periodic table

What makes people share? The degree to which a given bond is ionic or
covalent is determined by calculating the difference
Have you ever spent time with someone you really didn’t like? You had in electronegativity between the two atoms
nothing in common with them and did not want to have anything to do involved in the bond.
with them. On the other hand, there are people you enjoy being with.
You have a lot in common and like to share with them. Atoms work the As an example, consider the bond that occurs
same way. If there are strong differences in their attraction of electrons, between an atom of potassium and an atom of
one atoms gets the electrons and the other atom loses them. If they are fluorine. Using the table, the difference in
similar, they share the electrons to form a covalent bond. electronegativity is equal to 4.0 - 0.8 = 3.2. Since
the difference in electronegativity is relatively large,
Bond Polarity the bond between the two atoms is ionic. Since the
fluorine atom has a much larger attraction for
Electronegativity is defined as the ability of an atom to attract electrons electrons than the potassium atom does, the
when the atoms are in a compound. Electronegativities of elements are valence electron from the potassium atom is
shown in the periodic table below. completely transferred to the fluorine atom. The
diagram below shows how difference in
electronegativity relates to the ionic or covalent
character of a chemical bond.

BPaotntderPnoslanridtytrends in the periodic table

Non-polar Covalent Bonds Polar Covalent Bonds

A bond in which the electronegativity difference is less than 1.7 is A bond in which the electronegativity difference between the
considered to be mostly covalent in character. However, at this point we atoms is between 0.4 and 1.7 is called a polar covalent bond. A
need to distinguish between two general types of covalent bonds. A polar covalent bond is a covalent bond in which the atoms have
non-polar covalent bond is a covalent bond in which the bonding an unequal attraction for electrons and so the sharing is
electrons are shared equally between the two atoms. In a non-polar unequal. In a polar covalent bond, sometimes simply called a
covalent bond, the distribution of electrical charge is balanced between polar bond, the distribution of electrons around the molecule is
the two atoms. no longer symmetrical.

An easy way to illustrate the uneven electron distribution in a
polar covalent bond is to use the Greek letter delta ().

he two chlorine atoms share the pair of electrons in the single covalent The atom with the greater electronegativity acquires a partial
bond equally, and the electron density surrounding the Cl2 molecule is negative charge, while the atom with the lesser
symmetrical. Also note that molecules in which the electronegativity electronegativity acquires a partial positive charge. The delta
difference is very small (<0.4) are also considered non-polar covalent. An symbol is used to indicate that the quantity of charge is less
example would be a bond between chlorine and bromine (EN = 3.0-2.8 than one. A crossed arrow can also be used to indicate the
= 0.2) direction of greater electron density.

Summary

• The electronegativity of an atom determines how strongly it attracts electrons to itself.
• The polarity of a bond is affected by the electronegativity values of the two atoms involved in that bond.

Final wrap-up

Review
1. f two atoms bonded together have an electronegativity difference of 1.9, what is the bond

type?
2. What would be the bond type for BH2?
3. Your friend tells you that the LiF bond is covalent. Are they correct? Why or why not.

PaotlaterrMnsoalencdulteresnds in the periodic table Polar molecules orient themselves in the presence of an electric field
with the positive ends of the molecules being attracted to the negative
How cold are ultracold polar molecules? plate while the negative ends of the molecules are attracted to the
positive plate (see Figure below).
Ultracold polar molecules have been created by research
scientists. These “molecules” consists of Rb and K atoms Polar molecules are randomly oriented in the absence of an applied electric field
excited by lasers to form a type of Rb-K compound where (top). In an electric field, the molecules orient themselves because of the
the Rb has a positive charge and the K has a negative attraction of opposite charges (bottom).
charge. The material is formed at temperatures extremely
close to absolute zero. The researchers believe these
techniques will help them make new reactions and new
materials.

Polar Molecules

A polar molecule is a molecule in which one end of the
molecule is slightly positive, while the other end is slightly
negative. An example of a diatomic molecule that consists
of a polar covalent bond is hydrogen fluoride (HF). In
hydrogen fluoride, the hydrogen has a positive partial
charge and the fluorine has a negative partial charge. The
two electrically charged regions on either end of the
molecule are called poles, similar to a magnet having a
north and a south pole. A polar molecule is called a dipole
because it has two poles. Similarly, a polar bond itself can
be called a dipole. A dipole is indicated by an arrow that
starts on the positive end of the molecule and points to the
negative end (see Figure below).

PaotlaterrMnsoalencdulteresnds in the periodic table

For molecules with more than two atoms, the molecular geometry must Some other molecules are shown in the Figure below. Notice that a tetrahedral
also be taken into account when determining if the molecule is polar or
nonpolar. The Figure below shows a comparison between carbon dioxide molecule such as CH4 is nonpolar. However, if one of the peripheral H atoms is
and water. Carbon dioxide (CO2) is a linear molecule. The oxygen atoms replaced with another atom that has a different electronegativity, the molecule
are more electronegative than the carbon atom, so there are two
individual dipoles pointing outward from the C atom to each O atom. becomes polar. A trigonal planar molecule (BF3) may be nonpolar if all three
However, since the dipoles are of equal strength and are oriented in this peripheral atoms are the same, but a trigonal pyramidal molecule (NH3) is polar.
way, they cancel each other out and the overall molecular polarity of CO2
is zero.

Water is a bent molecule because of the two lone pairs on the central
oxygen atom. The individual dipoles point from the H atoms toward the O
atom. Because of the shape, the dipoles do not cancel each other out
and the water molecule is polar. In the Figure below, the net dipole is
shown in blue and points upward.

Self-learning 3.5
min

Before you join me in the classroom, make sure that
you watch this fun VDO beforehand.

Summary

• Polar molecules result from differences in electronegativity of the atoms in the molecule.
• Dipoles that are directly opposite one another cancel each other out.

Final wrap-up

Review
1. Why is a polar molecule called a dipole?
2. The molecular geometry of NH3 is trigonal pyramidal. What about the molecular

geometry of NH3 makes it a polar molecule?
3. What happens to polar molecules when an electric field is applied?

VPatletenrcnesBaonnddtrTehnedosryin the periodic table

What happens next?

We have seen that the old fish-hook idea of atoms connecting that Democritus liked so
much just doesn’t work. Electrons don’t have little hooks on them, but they are the basis
for connecting atoms to form molecules. You have learned how to write Lewis electron-dot
structures for molecules and predict their shape using VSEPR theory. Now it is time to apply
these abilities to understand how the electrons behave in their atomic orbitals when a
covalent bond forms.

Valence Bond Theory In essence, any covalent bond results from the overlap of atomic
orbitals. This idea forms the basis for a quantum mechanical
You have learned that a covalent bond forms when the electron clouds of two atoms theory called valence bond (VB) theory. In valence bond theory,
overlap with each other. In a simple H2 molecule, the single electron in each atom becomes the electrons in a molecule are assumed to occupy atomic
attracted to the nucleus of the other atom in the molecule as the atoms come closer orbitals of the individual atoms and a bond results from overlap
together. An optimum distance, equal to the bond length, is eventually attained, and the of those orbitals.
potential energy reaches a minimum. A stable, single covalent bond has formed between
the two hydrogen atoms. Other covalent bonds form in the same way as unpaired electrons
from two atoms “match up” to form the bond. In a fluorine atom, there is an unpaired
electron in one of the 2p orbitals. When a F2 molecule forms, the 2p orbitals from each of
the two atoms overlap to produce the F−F covalent bond. The overlapping orbitals do not
have to be of the same type. In a molecule of HF, the 1s orbital of the hydrogen atom
overlaps with the 2p orbital of the fluorine atom (see Figure on the right).

Summary

• Electrons occupy atomic orbitals.
• Covalent bonds result from the overlap of atomic orbitals.

Final wrap-up

Review
1. Where are electrons according to valence bond theory?
2. How do covalent bonds form?
3. Do the orbitals of the two electrons involved in the bond need to be the same?

HPaytbtreirdnOs rabnidtatlrse-nsdps3in the periodic table

Hybrid Orbitals – sp3 Now, four bonds are possible. The promotion of the electron “costs” a small
amount of energy, but recall that the process of bond formation is accompanied
The bonding scheme described by valence bond theory must account by a decrease in energy. The two extra bonds that can now be formed results in a
for molecular geometries as predicted by VSEPR theory. To do that, we lower overall energy and thus greater stability to the CH4 molecule. Carbon
must introduce a concept called hybrid orbitals. normally forms four bonds in most of its compounds.

sp3 Hybridization The number of bonds is now correct, but the geometry is wrong. The three p
orbitals (px, py, pz) are oriented at 90o relative to one another. However, as was
Unfortunately, overlap of existing atomic orbitals (s, p, etc.) is not seen from VSEPR theory, the observed H−C−H bond angle in the tetrahedral CH4
sufficient to explain some of the bonding and molecular geometries that molecule is actually 109.5o. Therefore, the methane molecule cannot be
are observed. Consider the element carbon and the methane (CH4) adequately represented by simple overlap of the 2s and 2p orbitals of carbon
molecule. A carbon atom has the electron configuration of 1s2 2s2 2p2, with the 1s orbitals of each hydrogen atom.
meaning that it has two unpaired electrons in its 2p orbitals, as shown
in Figure below. To explain the bonding in methane, it is necessary to introduce the concept of
hybridization and hybrid atomic orbitals. Hybridization is the mixing of the
Orbital configuration for carbon atom. atomic orbitals in an atom to produce a set of hybrid orbitals. When
hybridization occurs, it must do so as a result of the mixing of nonequivalent
According to the description of valence bond theory so far, carbon orbitals. In other words, s and p orbitals can hybridize but p orbitals cannot
would be expected to form only two bonds, corresponding to its two hybridize with other p orbitals. Hybrid orbitals are the atomic orbitals obtained
unpaired electrons. However, methane is a common and stable when two or more nonequivalent orbitals form the same atom combine in
molecule, with four equivalent C−H bonds. To account for this, one of preparation for bond formation. In the current case of carbon, the single 2s
the 2s electrons is promoted to the empty 2p orbital (see Figure below). orbital hybridizes with the three 2p orbitals to form a set of four hybrid orbitals,
called sp3 hybrids (see Figure below).
Promotion of carbon s electron to empty p orbital.
Carbon sp3 hybrid orbitals.

HPaytbtreirdnOs rabnidtatlrse-nsdps3in the periodic table

The sp3 hybrids are all equivalent to one another. Spatially, the
hybrid orbitals point towards the four corners of a tetrahedron
(see Figure on the right).

The process of sp3 hybridization is the mixing of an s orbital with a set of three p orbitals to form a
set of four sp3 hybrid orbitals. Each large lobe of the hybrid orbitals points to one corner of a
tetrahedron. The four lobes of each of the sp3 hybrid orbitals then overlap with the normal
unhybridized 1s orbitals of each hydrogen atoms to form the tetrahedral methane molecule.

Summary

• Electrons hybridize in order to form covalent bonds.
• Nonequivalent orbitals mix to form hybrid orbitals.

Final wrap-up

Review
1. Why is carbon expected to form only two covalent bonds?
2. How many covalent bonds does carbon actually form?
3. What needs to happen to allow carbon to form four bonds?

HPaytbtreirdnOs rabnidtatlrse-nsdpsaindthsep2periodic table

sp Hybridization The geometry of the sp hybrid orbitals is linear, with the lobes of the orbitals
pointing in opposite directions along one axis, arbitrarily defined as the x-axis
A beryllium hydride (BeH2) molecule is predicted to be linear by VSEPR. The (see Figure on the right). Each can bond with a 1s orbital from a hydrogen
beryllium atom contains all paired electrons and so must also undergo atom to form the linear BeH2 molecule

hybridization. One of the 2s electrons is first promoted to the empty 2px orbital
(see Figure below).

Promotion of Be 2s electron.

Now the hybridization takes place only with the occupied orbitals and the result is
a pair of sp hybrid orbitals. The two remaining p orbitals (py and pz) do not
hybridize and remain unoccupied (see Figure below).

Be hybrid orbitals. Other molecules whose electron domain geometry is linear and for whom

. hybridization is necessary also form sp hybrid orbitals. Examples include CO2
and C2H2, which will be discussed in further detail later.

PHaytbtreirdnOs rabnidtatlrse-nsdpsaindthsep2periodic table

sp2 Hybridization

Boron trifluoride (BF3) is predicted to have a trigonal planar geometry by VSEPR.
First a paired 2s electron is promoted to the empty 2py orbital (see Figure below).

Promotion of 2s electron.

This is followed by hybridization of the three occupied orbitals to form a set of
three sp2 hybrids, leaving the 2pz orbital unhybridized (see Figure below).

Formation of sp2 orbital. Other molecules with a trigonal planar electron domain geometry form sp2

The geometry of the sp2 hybrid orbitals is trigonal planar, with the lobes of the hybrid orbitals. Ozone (O3) is an example of a molecule whose electron
orbitals pointing towards the corners of a triangle (see Figure on the right). The domain geometry is trigonal planar, though the presence of a lone pair on
angle between any two of the hybrid orbital lobes is 120°. Each can bond with a
2p orbital from a fluorine atom to form the trigonal planar BF3 molecule. the central oxygen makes the molecular geometry bent. The hybridization of
the central O atom of ozone is sp2.

Summary

• Paired electrons can be hybridized and then participate in covalent bonding.

Final wrap-up

Review
1. Does the ground state beryllium atom contain any unpaired electrons?
2. Why does one 2s electron in Be get promoted to a 2p orbital?
3. What is the geometry of the two sp orbitals?

SPiagtmtearnasnadnPdi tBroends in the periodic table

Sigma and Pi Bonds The three sp2 hybrid orbitals lie in one plane, while the unhybridized 2pz orbital is
oriented perpendicular to that plane. The bonding in C2H4 is explained as follows. One
The hybridization model helps explain molecules with double or triple of the three sp2 hybrids forms a bond by overlapping with the identical hybrid orbital on
bonds (see Figure below). Ethene (C2H4) contains a double covalent bond
between the two carbon atoms and single bonds between the carbon the other carbon atom. The remaining two hybrid orbitals form bonds by overlapping
atoms and the hydrogen atoms. The entire molecule is planar.
with the 1s orbital of a hydrogen atom. Finally, the 2pz orbitals on each carbon atom
Geometry of ethene molecule. form another bond by overlapping with one another sideways.

As can be seen in Figure above, the electron domain geometry around It is necessary to distinguish Sigma and pi bonds.
each carbon independently is trigonal planar. This corresponds to sp2 between the two types of covalent
hybridization. Previously, we saw carbon undergo sp3 hybridization in a bonds in a C2H4 molecule. A sigma
CH4 molecule, so the electron promotion is the same for ethene, but the bond (σ bond) is a bond formed by
hybridization occurs only between the single s orbital and two of the the overlap of orbitals in an end-to-
three p orbitals. Thus generates a set of three sp2 hybrids along with an end fashion, with the electron
unhybridized 2pz orbital. Each contains one electron and so is capable of density concentrated between the
forming a covalent bond. nuclei of the bonding atoms. A pi
bond (π bond) is a bond formed by
Hybridization in ethene. the overlap of orbitals in a side-by-
side fashion with the electron
density concentrated above and
below the plane of the nuclei of the
bonding atoms.

Figure above shows the two types of bonding in C2H4. The sp2 hybrid orbitals are purple
and the pz orbital is blue. Three sigma bonds are formed from each carbon atom for a
total of six sigma bonds total in the molecule. The pi bond is the “second” bond of the

double bonds between the carbon atoms and is shown as an elongated green lobe that

extends both above and below the plane of the molecule. This plane contains the six

atoms and all of the sigma bonds.

SPiagtmtearnasnadnPdi tBroends in the periodic table

In a conventional Lewis electron-dot structure, a double bond is shown as The sp hybrid orbitals form a sigma bond between each other as well as sigma bonds to
a double dash between the atoms as in C=C. It is important to realize, the hydrogen atoms. Both the py and the pz orbitals on each carbon atom form pi
however, that the two bonds are different: one is a sigma bond, while the bonds between each other. As with ethene, these side-to-side overlaps are above and
other is a pi bond. below the plane of the molecule. The orientation of the two pi bonds is that they are
perpendicular to one another (see Figure below). One pi bond is above and below the
Ethyne (C2H2) is a linear molecule with a triple bond between the two line of the molecule as shown, while the other is in front of and behind the page.
carbon atoms (see Figure below). The hybridization is therefore sp.

Ethyne structure. The C2H2 molecule contains a triple bond between the two carbon atoms, one of which is a sigma
bond, and two of which are pi bonds.
The promotion of an electron in the carbon atom occurs in the same way.
However, the hybridization now involves only the 2s orbital and the 2px In general, single bonds between atoms are always sigma bonds. Double bonds are
orbital, leaving the 2py and the 2pz orbitals unhybridized. comprised of one sigma and one pi bond. Triple bonds are comprised of one sigma
bond and two pi bonds.

Hybridization in ethyne.

Summary

• Sigma bonds form between two atoms.
• Pi bonds form from p orbital overlap.

Final wrap-up

Review
1. What is the hybridization around each carbon in ethene?
2. What are the two types of bonds in C=C?
3. What is the shape of the ethene molecule?
4. How are the ethyne pi bonds oriented in relation to each

other?