Definition of HybridizationHybridization happens when
atomicorbitalsmix to form new atomic orbitals. The new orbitals
have the same totalelectroncapacity as the old ones. The properties
and energies of the new, hybridized orbitals are an 'average' of
the original unhybridized orbitals.The concept of hybridization was
introduced because it was the best explanation for the fact that
all of the C - H bonds in molecules like methane were
identical.ExampleCarbon atoms naturally have electron configuration
1s22s22p2.The four outermost electrons, i.e. those in the 2s and 2p
sublevels are available to form chemical bonds with other atoms.The
2s orbital is capable of holding up to two electrons, and there are
three 2p orbitals, each capable of holding up to two electrons,
which means the 2p orbitals can hold up to six
electrons.Individually, these electron orbitals look something like
this. (Each is centered on carbon's nucleus and the p orbitals make
angles of 90 with one another.)
The 2s orbital and the three 2p orbitals hybridize to form a new
hybrid orbital, called sp3, which is capable of holding up to eight
electrons.The sp3hybrid orbital, which is the sum of the original
2s and 2p orbitals, looks something like this:
sp3hybrid orbitals have a tetrahedral shape - each orbital makes
an angle of 109.5 with the others. This angle maximizes the
distance between the orbitals, which is natural given that the
electrons in each orbital repel one another.The shape of molecules
like methane, CH4, with bond angles of 109.5, is consistent with
sp3hybridization of carbon atoms.
Here is an energy level diagram showing how electron energies
change in hybridization.
Orbital hybridisationFrom Wikipedia, the free encyclopedia
Four sp3orbitals.
Three sp2orbitals.Inchemistry,hybridisation(orhybridization) is
the concept of mixingatomic orbitalsinto newhybrid orbitals(with
different energies, shapes, etc., than the actual orbitals
hybridising) suitable for the pairing of electrons to formchemical
bondsinvalence bond theory. Hybrid orbitals are very useful in the
explanation ofmolecular geometryand atomic bonding properties.
Although sometimes taught together with thevalence shell
electron-pair repulsion (VSEPR) theory, valence bond and
hybridisation are in fact not related to the VSEPR
model.[1]Contents[hide] 1Historical development 2spxand
sdxterminology 3Types of hybridisation 3.1sp3hybrids 3.2sp2hybrids
3.3sp hybrids 4Hybridisation and molecule shape 4.1Main group
compounds with lone pairs 5Hybridisation of hypervalent molecules
5.1Traditional description 5.2Coulson-Fischer description 5.2.1Main
group compounds with lone pairs 5.2.2Pi bonding 6Clarifying
misconceptions 6.1VSEPR electron domains and hybrid orbitals are
different 6.2Non-inclusion of dorbitals in main group compounds
6.3Non-inclusion of porbitals in transition metal complexes
7Hybridisation theory vs. MO theory 8See also 9References
10External links
Historical development[edit]ChemistLinus Paulingfirst developed
the hybridisation theory in order to explain the structure
ofmoleculessuch asmethane(CH4).[2]This concept was developed for
such simple chemical systems, but the approach was later applied
more widely, and today it is considered an effective heuristic for
rationalizing the structures oforganic compounds.Orbitals are a
model representation of the behaviour of electrons within
molecules. In the case of simple hybridisation, this approximation
is based on atomic orbitals, similar to those obtained for the
hydrogen atom, the only atom for which an exact analytic solution
to itsSchrdinger equationis known. In heavier atoms, like carbon,
nitrogen, and oxygen, the atomic orbitals used are the 2s and 2p
orbitals, similar to excited state orbitals for hydrogen. Hybrid
orbitals are assumed to be mixtures of these atomic orbitals,
superimposed on each other in various proportions. It provides
aquantum mechanicalinsight to Lewis structures. Hybridisation
theory finds its use mainly in organic chemistry.spxand
sdxterminology[edit]This terminology describes the weight of the
respective components of a hybrid orbital. For example, in methane,
the C hybrid orbital which forms eachcarbonhydrogenbond consists of
25% s character and 75% p character and is thus described as
sp3(read ass-p-three) hybridised.Quantum mechanicsdescribes this
hybrid as an sp3wavefunctionof the form N(s+3p), where N is
anormalization constant(here 1/2) and p is a p orbital directed
along the C-H axis to form asigma bond. The ratio of coefficients
(denoted in general) is3in this example. Since theelectron
densityassociated with an orbital is proportional to the square of
the wavefunction, the ratio of p-character to s-character is 2= 3.
The p character or the weight of the p component is N22= 3/4.For
atoms forming equivalent hybrids with no lone pairs, there is a
correspondence to the number and type of orbitals used. For
example, sp3hybrids are formed from one s and three p orbitals.
However, in all other cases, there is no such correspondence. The
two bond-forming hybrid orbitals of oxygen in water can be
described as sp4, which means that they have 20% s character and
80% p character, but doesnotimply that they are formed from one s
and four p orbitals. As a result, the amount ofp-character is not
restricted to integer values; i.e., hybridisations likesp2.5are
also readily described.In general, any two hybrid orbitals on the
same atom must be mutuallyorthogonal. For an atom with s and p
orbitals forming hybrids hiand hjwith included angle, the
orthogonality condition implies the relation:. The p-to-s ratio for
hybrid i is, and for hybrid j it is. The bond directed towards a
moreelectronegativesubstituent tends to have higherp-character as
stated inBent's rule. In the special case of equivalent hybrids on
the same atom, again with included angle, the equation reduces to
just. For example, BH3has a trigonal planar geometry, three 120
bond angles, three equivalent hybrids about the boron atom, and
thusbecomes, givingfor the p-to-s ratio. In other words,
sp2hybrids.An analogous notation is used to describe sdxhybrids.
For example, thepermanganate ion(MnO4-) has sd3hybridisation with
orbitals that are 25% s and 75% d.Types of
hybridisation[edit]sp3hybrids[edit]Hybridisation describes the
bonding atoms from an atom's point of view. That is, for a
tetrahedrally coordinated carbon (e.g.,methaneCH4), the carbon
should have 4 orbitals with the correct symmetry to bond to the 4
hydrogen atoms.Carbon'sground stateconfiguration is
1s22s22px12py1or more easily read:C
1s2s2px2py2pz
The carbon atom can utilize its two singly occupied p-type
orbitals (the designations pxpyor pzare meaningless at this point,
as they do not fill in any particular order), to form twocovalent
bondswith two hydrogen atoms, yielding the "free
radical"methyleneCH2, the simplest of thecarbenes. The carbon atom
can also bond to four hydrogen atoms by an excitation of an
electron from the doubly occupied 2s orbital to the empty 2p
orbital, so that there are four singly occupied orbitals.C*
1s2s2px2py2pz
As the additional bond energy more than compensates for the
excitation, the formation of four C-H bonds is energetically
favoured.Quantum mechanically, the lowest energy is obtained if the
four bonds are equivalent which requires that they be formed from
equivalent orbitals on the carbon. To achieve this equivalence,
theangular distributionsof the orbitals change via a linear
combination of the valence-shell (Core orbitals are almost never
involved in bonding) s and p wave functions[3]to form
foursp3hybrids.C*
1ssp3sp3sp3sp3
In CH4, four sp3hybrid orbitals are overlapped byhydrogen's 1s
orbital, yielding four (sigma) bonds(that is, four single covalent
bonds) of the same length and strength.translates
intosp2hybrids[edit]
Ethene structureOther carbon based compounds and other molecules
may be explained in a similar way as methane. Take, for
example,ethene(C2H4). Ethene has a double bond between the
carbons.For this molecule, carbon will sp2hybridise, because one
(pi) bondis required for thedouble bondbetween the carbons, and
only three bonds are formed per carbon atom. Insp2hybridisationthe
2s orbital is mixed with only two of the three available 2p
orbitals:C*
1ssp2sp2sp22p
forming a total of three sp2orbitals with one p orbital
remaining. In ethylene (ethene) the two carbon atoms form a bond by
overlapping two sp2orbitals and each carbon atom forms two covalent
bonds with hydrogen by ssp2overlap all with 120 angles. The bond
between the carbon atoms perpendicular to the molecular plane is
formed by 2p2p overlap. The hydrogencarbon bonds are all of equal
strength and length, which agrees with experimental data.sp
hybrids[edit]
A schematic presentation of hybrid orbitals spThe chemical
bonding in compounds such asalkyneswith triple bonds is explained
bysp hybridisation.C*
1sspsp2p2p
In this model, the 2sorbital mixes with only one of the three
porbitals resulting in two sporbitals and two remaining unchanged
porbitals. The chemical bonding inacetylene(ethyne) (C2H2) consists
of spsp overlap between the two carbon atoms forming a bond and two
additionalbondsformed by pp overlap. Each carbon also bonds to
hydrogen in a ssp overlap at 180 angles.Hybridisation and molecule
shape[edit]Hybridisation helps to explainmolecule
shape:ClassificationMain groupTransition metal[4]
AX2 Linear(180) sp hybridisation E.g., CO2 Bent(90) sd
hybridisation E.g., VO2+
AX3 Trigonal planar(120) sp2hybridisation E.g., BCl3 Trigonal
pyramidal(90) sd2hybridisation E.g., CrO3
AX4 Tetrahedral(109.5)
sp3hybridisation E.g., CCl4 sd3hybridisation E.g., MnO4
AX5- Square pyramidal(66, 114)[5] sd4hybridisation E.g.,
Ta(CH3)5
AX6- Trigonal prismatic(63, 117)[5] sd5hybridisation E.g.,
W(CH3)6
Main group compounds with lone pairs[edit]For main group
compounds with lone electron pairs, the s orbital lone pair
(analogous to s-p mixing inmolecular orbital theory) can be
hybridised to a certain extent with the bond pairs[6]to maximize
energetic stability according to itsWalsh diagram. This
rationalisation is applied to explain deviations in ideal bond
angles (i.e. only p orbitals used for bonding), most commonly in
second and third period elements. Trigonal pyramidal(AX3E1)
s-orbital can be hybridised with the three p-orbital bonds to give
bond angles greater than 90. E.g., NH3 Bent(AX2E1-2) s-orbital lone
pair can be hybridised with the two p-orbital bonds to give bond
angles greater than 90. The out-of-plane p-orbital can either be a
lone pair or pi bond. If it is a lone pair, it results in
inequivalent lone pairs contrary to the common picture depicted by
VSEPR theory (see below). E.g., SO2, H2O Monocoordinate (AX1E1-3)
s-orbital lone pair can be hybridised with the p-orbital bond. The
two out-of-line p-orbitals can either be lone pairs or pi bonds.
The p-orbital lone pairs are inequivalent from the s-rich lone
pair. E.g., CO, SO, HFHybridisation of hypervalent
molecules[edit]Traditional description[edit]In general chemistry
courses and mainstream textbooks, hybridisation is often presented
for main group AX5 and above, as well as for transition metal
complexes, using the hybridisation scheme first proposed by
Pauling.ClassificationMain groupTransition metal
AX2- Linear(180) sp hybridisation E.g., Ag(NH3)2+
AX3- Trigonal planar(120) sp2hybridisation E.g., Cu(CN)32
AX4- Square planar(90) sp2d hybridisation E.g., PtCl42
AX5 Trigonal bipyramidal(90, 120) sp3d hybridisation E.g., PCl5,
Fe(CO)5
AX6 Octahedral(90) sp3d2hybridisation E.g., SF6, Mo(CO)6
AX7 Pentagonal bipyramidal(90, 72) sp3d3hybridisation E.g., IF7,
V(CN)74
AX8 Square antiprismatic sp3d4hybridisation E.g., IF8,
Re(CN)83
AX9- Tricapped trigonal prismatic sp3d5hybridisation E.g.,
ReH92
However, such a scheme is now superseded as more recent
calculations based on molecular orbital theory have shown that in
main-group molecules the d component is insignificant, while in
transition metal complexes the p component is insignificant (see
below).Coulson-Fischer description[edit]In
theCoulson-Fischerdescription, the singly-occupied bond orbitals
utilize a ligand orbital component to account for the ionicity of
the bond.[7]As shown by computational chemistry,hypervalent
moleculescan only be stable given strongly polar (and weakened)
bonds with electronegative ligands such as fluorine or oxygen to
reduce the valence electron occupancy of the central atom to below
8 (or 12 for transition metals). The added ligand orbital
component, which translates to ionic character, also allows the
central atom to form more orbitals for bonding than would be
possible with non-polar covalent bonds. This orbital picture can be
illustrated in Lewis structures by
ionic-covalentresonance.ClassificationMain groupTransition
metal[4]
AX2- Linear(180) A(s)+X() E.g., Ag(NH3)2+
AX3- Trigonal planar(120) A(s)+X() E.g., Cu(CN)32
AX4- Square planar(90) A(sd)+X() E.g., PtCl42
AX5 Trigonal bipyramidal(90, 120)
A(sp3)+X() E.g., PCl5 A(sd)+X() E.g., Fe(CO)5
AX6 Octahedral(90)
A(sp3)+X() E.g., SF6 A(sd2)+X() E.g., Mo(CO)6
AX7 Pentagonal bipyramidal(90, 72)
A(sp3)+X() E.g., IF7 A(sd3)+X() E.g., V(CN)74[8]
AX8 Square antiprismatic
A(sp3)+X() E.g., IF8[9] A(sd4)+X() E.g., Re(CN)83[10]
AX9- Tricapped trigonal prismatic A(sd5)+X() E.g., ReH92
Main group compounds with lone pairs[edit]For hypervalent main
group compounds with lone electron pairs, the bonding scheme can be
split into two components: the "hypervalent bonding" component and
the "regular bonding" component. The "regular bonding" component
has the same description (see above), while the "hypervalent
bonding" component consists of A(p)+X() hybrids. The table below
shows how each shape is related to the two components and their
respective descriptions.Regular bonding component
BentMonocoordinate-
Hypervalent bonding component(A(p)+X())Linear axis Seesaw(AX4E1)
(90, 180, >90) - E.g., SF4 T-shaped(AX3E2) (90, 180) + p-orbital
lone pair E.g., ClF3 Linear(AX2E3) (180) + two p-orbital lone
pairs, s-orbital lone pair E.g., XeF2
Square planar equator- Square pyramidal(AX5E1) (90, 90) - E.g.,
ClF5 Square planar(AX4E2) (90) + p-orbital lone pair, s-orbital
lone pair E.g., XeF4
Pentagonal planar equator- Pentagonal pyramidal(AX6E1) (90, 72)
- E.g., XeOF5 Pentagonal planar(AX5E2) (72) + p-orbital lone pair,
s-orbital lone pair E.g., XeF5
Pi bonding[edit]Hypervalent pi bonding in the Coulson-Fischer
description is also explained analogously using A(p)+X(p) or
A(d)+X(p) hybrids. For example, in ozone, the central oxygen's
orbital hybridises with the two surrounding oxygen orbitals and
overlaps with them to form two polar pi bonds with bond order of
0.5 each.Clarifying misconceptions[edit]VSEPR electron domains and
hybrid orbitals are different[edit]The simplistic picture of
hybridisation taught in conjunction with VSEPR theory does not
agree with high-level theoretical calculations[6]despite its
widespread usage in many textbooks. For example, following the
guidelines of VSEPR, the hybridization of the oxygen in water is
described as sp3with two lone electron-pairs and two bonds in four
equal-energy, symmetrical orbitals making a
tetrahedron.[11]However,molecular orbitalcalculations give orbitals
which reflect the symmetry of the molecule.[12]One of the two lone
pairs is in a pure p-type orbital, with its electron density
perpendicular to the H-O-H framework.[12]The other lone pair is in
an approximately sp0.8orbital that is in the same plane as the
H-O-H bonding.[12]Photoelectron spectraconfirm the presence of two
different energies for the nonbonded electrons.[13]In contrast, the
orbitals used to make the O-H bonds are approximately
sp4hybrids.[14]It has been argued that it is this change in the
mixing of the orbitals that is responsible for the compression of
the H-O-H angle down to the experimental 104.5degrees, not some
change in the repulsion of electrons.[14]Thus while VSEPR and its
application to hybridisation predicts the correct atomic framework
for water and other molecules with lone electron pairs, it may do
so for the wrong reason.Non-inclusion of dorbitals in main group
compounds[edit]Main article:Hypervalent moleculeIn 1990, Magnusson
published a seminal work definitively excluding the role of
d-orbital hybridization in bonding in hypervalent compounds of
second-row elements. This had long been a point of contention and
confusion in describing these molecules using molecular orbital
theory. Part of the confusion here originates from the fact that
one must include d-functions in the basis sets used to describe
these compounds (or else unreasonably high energies and distorted
geometries result), and the contribution of the d-function to the
molecular wavefunction is large. These facts were historically
interpreted to mean that d-orbitals must be involved in bonding.
However, Magnusson concludes in his work that d-orbital involvement
is not implicated in hypervalency.[15]Non-inclusion of porbitals in
transition metal complexes[edit]Similarly, porbitals have long been
thought to be utilized by transition metal centers in bonding with
ligands, hence the18-electrondescription; however, recentmolecular
orbitalcalculations have found that porbitals do not contribute
significantly to the hybrid orbitals in transition metal
complexes,[16]even though the contribution of the p-function to the
molecular wavefunction is calculated to be somewhat larger than
that of the d-function in main group compounds.Hybridisation theory
vs. MO theory[edit]Hybridisation theory is an integral part
oforganic chemistryand in general discussed together withmolecular
orbital theoryin advanced organic chemistry textbooks although for
different reasons. One textbook notes that for drawing reaction
mechanisms sometimes a classical bonding picture is needed with two
atoms sharing two electrons.[17]It also comments that predicting
bond angles in methane with MO theory is not straightforward.
Another textbook treats hybridisation theory when explaining
bonding in alkenes[18]and a third[19]uses MO theory to explain
bonding in hydrogen but hybridisation theory for methane.Bonding
orbitals formed from hybrid atomic orbitals may be considered
aslocalized molecular orbitals, which can be formed from the
delocalized orbitals of molecular orbital theory by an appropriate
mathematical transformation. For molecules with a closed electron
shell in the ground state, this transformation of the orbitals
leaves the total many-electron wave function unchanged. The hybrid
orbital description of the ground state is thereforeequivalentto
the delocalized orbital description for explaining the ground state
total energy and electron density, as well as the molecular
geometry which corresponds to the minimum value of the total
energy.There is no such equivalence, however, for ionized or
excited states with open electron shells. Hybrid orbitals cannot
therefore be used to interpret photoelectron spectra, which measure
the energies of ionized states, identified with delocalized orbital
energies usingKoopmans' theorem. Nor can they be used to interpret
UV-visible spectra which correspond to electronic transitions
between delocalized orbitals. From a pedagogical perspective,
hybridisation approach tends to over-emphasize localisation of
bonding electrons and does not effectively embracemolecular
symmetryas does MO theory.
IntroductionBased on Valence Bond Theory, carbon would only be
able to form two covalent bonds, making CH2. However, as you will
find out, we know that this is not true and that in reality, it
makes CH4. Thehybridizationof orbitals is also greatly favored
because hybridized orbitals are lower in energy compared to their
separated, unhybridized counterparts. This results in more stable
compounds whenhybridizationoccurs. Also, major parts of the
hybridized orbitals, or the frontal lobes, overlap better than the
lobes of unhybridized orbitals. This leads to better bonding.Carbon
is a perfect example showing the need for hybrid orbitals. As you
know, Carbon's ground state configuration is:
According to the valence bond theory, carbon should form two
covalent bonds, resulting in a CH2. However, tests show that CH2is
highly reactive and cannot exist outside of a reaction. Therefore,
this does not explain how CH4can exist. However, you can excite a
2s electron and bump it into one of the 2p orbitals. This would
give you the following configuration:
While this would allow us to have four covalent bonds, resulting
in CH4, it also implies that the C-H covalent bonds would have
different energies due to the different levels of orbital overlap.
However, with testing, it has been proven that in CH4, any hydrogen
can be removed with the same amount of energy. This means that
every C-H covalent bond should have equal energies. Once again,
this means that the valence bond theory fails to explain the
existence of CH4. The only way it can be explained is if when we
had the exited state above, the 2s and the 3 2p orbitals fused
together to make four, equal energysp3hybrid orbitals. That would
give us the following configuration:
This explains how a carbon can have four equal energy bonds. The
next section will explain the various types ofhybridizationand how
each type helps explain the structure of certain
molecules.spHybridizationspHybridizationcan explain the linear
structure in molecules. In it, the 2s orbital and one of the 2p
orbitals hybridize to form two sp orbitals, each consisting of 50%
s and 50% p character. The front lobes face away from each other
and form a straight line leaving a 180 angle between the two
orbitals. This formation minimizes electron repulsion. Because only
one p orbital was used, we are left with two unaltered 2p orbitals
that the atom can use. These p orbitals are at right angles to one
another and to the line formed by the two sp orbitals.Energy
changes occurring inhybridization
Notice how the energy of the electrons lowers when
hybridized.These p orbitals come into play in compounds such as
ethyne where they form two addition ? bonds, resulting in in a
triple bond. This only happens when two atoms, such as two carbons,
both have two p orbitals that each contain an electron. An sp
hybrid orbital results when an s orbital is combined with p orbital
(Figure 2). We will get two sp hybrid orbitals since we started
with two orbitals (s and p). sphybridizationresults in a pair of
directional sp hybrid orbitals pointed in opposite directions.
These hybridized orbitals result in higher electron density in the
bonding region for a sigma bond toward the left of the atom and for
another sigma bond toward the right. In addition,
sphybridizationprovides linear geometry with a bond angle of
180o.Examples of spHybridizationMagnesium HydrideIn magnesium
hydride, the 3s orbital and one of the 3p orbitals from magnesium
hybridize to form two sp orbitals. The two frontal lobes of the sp
orbitals face away from each other forming a straight line leading
to a linear structure. These two sp orbitals bond with the two 1s
orbitals of the two hydrogen atoms through sp-s orbital
overlap.Hybridization
EthyneThehybridizationin ethyne is similar to thehybridizationin
magnesium hydride. For each carbon, the 2s orbital hybridizes with
one of the 2p orbitals to form two sp hybridized orbitals. The
frontal lobes of these orbitals face away from each other forming a
straight line. The first bond consists of sp-sp orbital overlap
between the two carbons. Another two bonds consist of s-sp orbital
overlap between the sp hybridized orbitals of the carbons and the
1s orbitals of the hydrogens. This leaves us with two p orbitals on
each carbon that have a single carbon in them. This allows for the
formation of two ? bonds through p-p orbital overlap. The linear
shape, or 180 angle, is formed because electron repulsion is
minimized the greatest in this position.Hybridization
sp2hybridizationsp2hybridizationcan explain the trigonal planar
structure of molecules. In it, the 2s orbitals and two of the 2p
orbitals hybridize to form three sp orbitals, each consisting of
67% p and 33% s character. The frontal lobes align themselves in
the trigonal planar structure, pointing to the corners of a
triangle in order to minimize electron repulsion and to improve
overlap. The remaining p orbital remains unchanged and is
perpendicular to the plane of the three sp2orbitals.
Energy changes occurring inhybridization
Hybridizationof an s orbital with two p orbitals (pxand py)
results in three sp2hybrid orbitals that are oriented at 120oangle
to each other (Figure 3). Sp2hybridizationresults in trigonal
geometry.Examples of sp2hybridizationAluminum TrihydrideIn aluminum
trihydride, one 2s orbital and two 2p orbitals hybridize to form
three sp2orbitals that align themselves in the trigonal planar
structure. The three Al sp2orbitals bond with with 1s orbitals from
the three hydrogens through sp2-s orbital overlap.Hybridization
EtheneSimilarhybridizationoccurs in each carbon of ethene. For
each carbon, one 2s orbital and two 2p orbitals hybridize to form
three sp2orbitals. These hybridized orbitals align themselves in
the trigonal planar structure. For each carbon, two of these sp
orbitals bond with two 1s hydrogen orbitals through s-sp orbital
overlap. The remaining sp2orbitals on each carbon are bonded with
each other, forming a bond between each carbon through
sp2-sp2orbital overlap. This leaves us with the two p orbitals on
each carbon that have a single carbon in them. These orbitals form
a ? bonds through p-p orbital overlap, creating a double bond
between the two carbons. Because a double bond was created, the
overall structure of the ethene compound is linear. However, the
structure of each molecule in ethene, the two carbons, is still
trigonal planar.
Hybridization
sp3hybridizationsp3hybridizationcan explain the tetrahedral
structure of molecules. In it, the 2s orbitals and all three of the
2p orbitals hybridize to form four sp orbitals, each consisting of
75% p character and 25% s character. The frontal lobes align
themselves in the manner shown below. In this structure, electron
repulsion is minimized.Energy changes occurring inhybridization
Hybridizationof an s orbital with all three p orbitals (px, py,
and pz) results in four sp3hybrid orbitals. sp3hybrid orbitals are
oriented at bond angle of 109.5ofrom each other. This
109.5oarrangement gives tetrahedral geometry (Figure 4).Examples of
sp3hybridizationMethaneBecause carbon plays such a significant role
in organicchemistry, we will be using it as an example here.
Carbon's 2s and all three of its 3p orbitals hybridize to form four
sp3orbitals. These orbitals then bond with four hydrogen atoms
through sp3-s orbital overlap, creating methane. The resulting
shape is tetrahedral, since that minimizes electron
repulsion.Hybridization
Lone PairsRemember to take into account lone pairs of electrons.
These lone pairs cannot double bond so they are placed in their own
hybrid orbital. This is why H2O is tetrahedral.We can also build
sp3d and sp3d2hybrid orbitals if we go beyond s and p
subshells.References1. John Olmsted, Gregory M. WilliamsChemistry:
The Molecular ScienceJones & Bartlett Publishers 1996.
366-3712. Francis A. CareyAdvanced OrganicChemistrySpringer 2001.
4-63. L. G. Wade, Jr. Whitman CollegeOrganicChemistryFifth Edition
2003Practice ProblemsUsing theLewis Structures, try to figure out
thehybridization(sp, sp2, sp3) of the indicated atom and indicate
the atom's shape.1. The carbon.
2. The oxygen.
3. The carbon on the right.
Answers1. sp2- Trigonal PlanarThe carbon has no lone pairs and
is bonded to three hydrogens so we just need three hybrid orbitals,
aka sp2.2. sp3- TetrahedralDon't forget to take into account all
the lone pairs. Every lone pair needs it own hybrid orbital. That
makes three hybrid orbitals for lone pairs and the oxygen is bonded
to one hydrogen which requires another sp3orbital. That makes 4
orbitals, aka sp3.3. sp - LinearThe carbon is bonded to two other
atoms, that means it needs two hybrid orbitals, aka sp.An easy way
to figure out whathybridizationan atom has is to just count the
number of atoms bonded to it and the number of lone pairs. Double
and triple bonds still count as being only bonded to one atom. Use
this method to go over the above problems again and make sure you
understand it. It's a lot easier to figure out thehybridizationthis
way.Contributors Harpreet Chima (UCD), Farah YasmeenThis pageviewed
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byMindtouchCore 2010Sigma andPi BondsA description of thedouble
bondis thesigma-pimodel shown in Fig. 1. In this case only two of
theporbitals on each Catomare involved in the formation of hybrids.
Consequentlysp2hybrids are formed, separated by an angle of 120.
Two of these hybrids from each C atom overlap with H 1sorbitals,
while the third overlaps with ansp2hybrid on the other C atom. This
overlap directly between the two C atoms is called asigma bond, and
is abbreviated by the Greek letter . This orbital has no
nodes:electrondensityexists continuously from around one atom to
the other atom.
Figure 1 The sigma-pi model of a double bond. Three sp2hybrids
around each carbon atom are indicated in color. Two of these
overlap directly between the carbon atoms to form the bond. Two p
orbitals, one on each C atom, are shown in gray. These overlap
sideways to form a bond, also shown in gray.
Top of FormBottom of FormTop of FormNo MOHOMOLUMOBottom of
FormTop of FormMO cutoff:Bottom of Form
To view the sigma bonding orbital, selectN6. This is actually
sigma bonding between C-C and some sigma-like bonding around the Hs
as well. Focus on the yellow portion. By selectingN8 HOMO, you can
see the pi orbital represented by the two lobes.
The sp2hybrid orbitals on each carbon atom involve the 2sand two
of the 2porbitals, leaving a single 2porbital on each carbon atom.
A second carbon-carbon bond is formed by the overlap of these two
remainingporbitals. This is called api bond, Greek letter . The pi
bond ( bond) has two halvesone above the plane of themolecule, and
the other below it. Each of the two electrons in the pi bond (
bond) exists both above and below the plane of the four H atoms and
the two C atoms. The pi bond can be thought of as a standing wave
with a single node in the plane of the molecule.If your workstation
is enabled for JCE Software, you will see two videos below which
compare the behavior of a standing wave with zero nodes versus a
standing wave with one node (otherwise, see the drum animation
below). The wave with a single node has higherenergy. The sigma
bond between the two carbon atoms does not have a node in the plane
of the molecule. The pi bond between the two carbon atoms has one
node in the plane of the molecule. Thus the pimolecular orbitalis
higher in energy and is the highest occupied molecular orbital (the
HOMO).
(a) No Nodes(b) Single NodeStanding Waves on a Wave
Demonstrator. (a) The lowest energy form of a standing wave has no
nodes. This is like the continual electron density in all
directions around the sigma bonding orbital. (b) The second lowest
energy standing wave has a single node. This node is akin to the
shape of the pi bond where there is no electron density along the
plane.
Alternatively, we can envision the molecular orbitals with
theDrum Modeldescribed earlier. Imagine the two atoms opposite one
another where a diagonal meets the edge of the drum at extreme left
and right points. The m1,0mode has no nodes, so the maximum
amplitude of the standing wave is between the atoms, representing a
high electron density sigma bond. The m2,1mode has a linear node
between the atoms, and maximum amplitude in front of, and behind,
the node, representing the pi bond. In 3D, this linear node would
be a plane, separating the two lobes of high electron density that
constitute the pi bond. Because the pi bond has less electron
density between the atoms, it is ofhigher energyin the MO diagram
and isweakerthan the sigma bond.mode m1,0mode m2,1
Overall this sigma-pi picture of the double bond is reminiscent
of a hot dog in a bun. The sigma bond ( bond) corresponds to the
frankfurter, while the pi bond corresponds to the bun on either
side of it.Although the sigma-pi picture is more complex than the
bent-bond picture of the double bond, it is much used
byorganicchemists (those chemists interested in carboncompounds).
The sigma-pi model is especially helpful in understanding what
happens when visible light or other radiation is absorbed by a
molecule. Further discussion on this topic is found in the sections
onSpectra and Structure of Atoms and Molecules.
Figure 2 Dot-density diagrams comparing the sigma-pi and
bent-bond models of the double bond. (a) The sigma-pi model. The
bond is in color and the bond is in gray. (b) The bent-bond model.
One lies above the two nuclei and the other below. Since both
diagrams are dot-for-dot the same, they are both describing the
same physical reality.In actual fact the difference between the two
models of the double bond (the first model described here and the
second found in the section onOrbital Descriptions of Multiple
Bonds) is more apparent than real. They are related to each other
in much the same way assandporbitals are related tosphybrids.
Figure 2 shows two dot-density diagrams for a carbon-carbon double
bond in a plane through both carbon nuclei but at right angles to
the plane of the molecule. Figure 2acorresponds to a sigma-pi model
with the sigma bond ( bond) in color and the pi bond in gray.
Figure 2bshows two bent bonds. Careful inspection reveals that both
diagrams are dot-for-dot the same. Only the color coding of the
dots is different. Thus the bent-bond and sigma-pi models of the
double bond are just two different ways of dividing up the same
overall electron density.A similar situation applies totriple
bonds, such as that found in a molecule of ethyne (acetylene),. As
shown in Fig. 3a,we can regard this triple bond as being the result
of three overlaps ofsp3hybrids on different carbon atoms forming
three bent bonds. Alternatively we can regard it as being composed
of one sigma bond and two pi bonds, the sigma bond being due to the
overlap of ansphybrid from each carbon atom. Again both pictures of
the bond correspond to the same overall electron density, and hence
both are describing the same physical reality. We can use whichever
one seems more convenient for the problem under consideration.
Figure 3 Two alternative models for the triple bond in ethyne,.
(a) Threesp3hybrids from each carbon atom overlap to form three
bent bonds. (b) Twosphybrids overlap to form the sigma bond.
Twoporbitals on one carbon overlap with two on the other to form
two pi bonds (one in light gray, the other in dark gray). Though
these two models appear to be different, the indistinguishability
of electrons makes them exactly equivalent.
Top of FormBottom of FormTop of FormNo MOHOMOLUMOBottom of
FormTop of FormMO cutoff:Bottom of Form
The orbitals can be viewed by selecting from the orbital menu at
right. Again, these orbitals are more easily represented using an
MO Cutoff of 0.005. By selecting By selectingN1you can see the
sigma bonding orbital. The two pi bonding orbitals can be viewed by
selectingN6andN7. If you rotate the molecule so that we view the
molecule along the H-C-C-H line, you can switch back and forth
betweenN6andN7to see the orientation in space.
Sigma and Pi Bonds
Simply put, a sigma bond is a single covalent bond.The electron
pair is located between the two atoms involved in the bonding.A pi
bond uses the p-orbitals that are located above and below these
atoms.
The overlap is a pi-bond. The image above is actually only 1
pi-bond.A p-orbital is has a shape of a dumbbell. So there are 2
regions of overlapping.
So, the grey bond is a sigma bond (a single bond), the clouds
are a pi (this is the second bond or your double bond).So, how can
we have triple bonds?Use the image belowThe region of space above
and below the sigma bond (single bond) are already occupied. The
p-orbitals (Pink) can wrap around to the left and right of the
sigma bond. This overlap is 90ofrom the other pi-bond (blue) that
is already in place. So it is possible to have 2-pi bonds and a
sigma or what we call a triple bond.
In conclusion, a triple bond is a sigma bond located directly
between the atoms, and 2 pi bonds located above and below, and
around the sides of the 2 atoms.Every bond has a sigma. Doubles
have a sigma and a pi. Triples have a sigma and two pi bonds.
