Hybridization and Resonance
Hybridization:
In chemistry, hybridisation (or hybridization) is the idea
of mixing atomic orbitals to form new hybrid orbitals suitable for the
qualitative description of atomic bonding properties. Hybridised orbitals are
very useful in the explanation of the shape of molecular orbitals for
molecules. It is an essential part of valence bond theory. Although sometimes
taught together with the valence shell electron-pair repulsion (VSEPR) theory,
valence bond and hybridization are in fact not related to the VSEPR model.
Types of Hybridizations
There are Three types of hybridization which are given below:
- sp3
hybrids
- sp2
hybrids
- sp
hybrids
sp3
hybrids:
Hybridisation describes the bonding atoms from an atom's
point of view. That is, for a tetrahedrally coordinated carbon (e.g., methane,
CH4), the carbon should have 4 orbitals with the correct symmetry to bond to
the 4 hydrogen atoms. The problem with the existence of methane is now this:
carbon's ground-state configuration is 1s2 2s2 2px1 2py1 or more easily read:
The valence bond theory would predict, based on the
existence of two half-filled p-type orbitals (the designations px py or pz are
meaningless at this point, as they do not fill in any particular order), that C
forms two covalent bonds, i.e., CH2 (methylene). However, methylene is a very
reactive molecule (see also: carbene) and cannot live outside of a molecular
system. Therefore, this theory alone cannot explain the survival of CH4.
Furthermore, ground state orbitals cannot be used for
bonding in CH4. While exciting a 2s electron into a 2p orbital would, in
theory, allow for four bonds according to the valence bond theory, (which has
been proved experimentally accurate for systems like O2) this would imply that
the various bonds of CH4 would have differing energies due to differing levels
of orbital overlap. Once again, this has been experimentally disproved: any
hydrogen can be removed from a carbon with equal ease.
To summarise, to explain the existence of CH4 (and many
other molecules) a method by which as many as 12 bonds (for transition metals)
of equal strength (and therefore equal length) was required.
The first step in hybridisation is the excitation of one (or
more) electrons (we consider the carbon atom in methane, for simplicity of the
discussion):
The proton that forms the nucleus of a hydrogen atom
attracts one of the lower-energy valence electrons on carbon. This causes an
excitation, moving a 2s electron into a 2p orbital. This, however, increases
the influence of the carbon nucleus on the valence electrons by increasing the
effective core potential (the amount of charge the nucleus exerts on a given
electron = Charge of Core − Charge of all electrons closer to the nucleus). The
effective core potential is also known as the effective nuclear charge, or
Z_eff.
The solution to the Schrödinger equation for this
configuration is a linear combination of the s and p wave functions, or
orbitals, known as a hybridized orbital[3]. In the case of carbon attempting to
bond with four hydrogens, four orbitals are required. Therefore, the 2s orbital
(core orbitals are almost never involved in bonding) "mixes" with the
three 2p orbitals to form four sp3 hybrids (read as s-p-three). See graphical
summary below.
In CH4, four sp3 hybridised orbitals are overlapped by
hydrogen's 1s orbital, yielding four σ (sigma) bonds (that is, four single
covalent bonds). The four bonds are of the same length and strength. This
theory fits our necessities.
If we now recombine these orbitals with the blank s-orbitals
of 4 hydrogens (4 protons, H+) and allow maximum separation between the 4
hydrogens (i.e., tetrahedral surrounding of the carbon), we see that at any
orientation of the p-orbitals, a single hydrogen has an overlap of 25% with the
s-orbital of the C, and a total of 75% of overlap with the 3 p-orbitals (see
that the relative percentages are the same as the character of the respective
orbital in an sp3-hybridisation model, 25% s- and 75% p-character).
According to the orbital hybridisation theory, the valence
electrons in methane should be equal in energy but its photoelectron spectrum
[4] shows two bands, one at 12.7 eV (one electron pair) and one at 23 eV (three
electron pairs). This visible inconsistency can be explained when one considers
additional orbital mixing taking place when the sp3 orbitals mix with the 4
hydrogen orbitals.
sp2
hybrids:
Other 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 sp2 hybridise, because
one π (pi) bond is required for the double bond between the carbons, and only
three σ bonds are formed per carbon atom. In sp2 hybridisation the 2s orbital
is mixed with only two of the three available 2p orbitals
forming a total of 3 sp2 orbitals with one p-orbital
remaining. In ethylene (ethene) the two carbon atoms shape a σ bond by
overlapping two sp2 orbitals and each carbon atom forms two covalent bonds with
hydrogen by s–sp2 overlap all with 120° angles. The π bond between the carbon
atoms perpendicular to the molecular plane is formed by 2p–2p overlap. The
hydrogen-carbon bonds are all of equal strength and length, which agrees with
experimental data.
The amount of p-character is not restricted to integer values;
i.e., hybridisations like sp2.5 are also willingly described. In this case the
geometries are somewhat distorted from the ideally hybridised picture. For
example, as stated in Bent's rule, a bond tends to have higher p-character when
directed toward a more electronegative substituen
sp3
hybrids:
The chemical bonding in compounds such as alkynes with
triple bonds is explained by sp hybridization.
In this model, the 2s orbital mixes with only one of the
three p-orbitals resulting in two sp orbitals and two remaining unchanged p
orbitals. The chemical bonding in acetylene (ethyne) (C2H2) consists of sp–sp
overlap between the two carbon atoms forming a σ bond and two additional π
bonds formed by p–p overlap. Each carbon also bonds to hydrogen in a sigma s–sp
overlap at 180° angles.
Resonance :
In chemistry, resonance or mesomerism
is a way of describing delocalized electrons
within certain molecules or polyatomic ions wherever the bonding cannot be
expressed by one single Lewis formula. A molecule or ion with such delocalized
electrons is represented by several contributing structures
(also called resonance structures or canonical
forms).
Every contributing structure can be represented by a Lewis
structure, with only an integer number of covalent bonds between each pair of
atoms within the structure. These individual contributors cannot be observed in
the actual resonance-stabilized molecule; resonance is not a
rapidly-interconverting set of contributors. Several Lewis structures are used
collectively to describe the actual molecular structure. The actual structure
is an approximate intermediate between the canonical forms, but its overall
energy is lower than each of the contributors. This intermediate form between
different contributing structures is called a resonance hybrid. Contributing
structures differ only in the position of electrons, not in the position of
nuclei.
Resonance is a key component of valence bond theory.
Electron delocalization lowers the potential energy of the
substance and thus makes it more stable than any of the contributing
structures. The difference between the potential energy of the real structure
and that of the contributing structure with the lowest potential energy is
called the resonance energy or delocalization energy.
General characteristics of resonance:
- They
can be represented by several correct Lewis formulas, called
"contributing structures", "resonance structures" or
"canonical forms". However, the real structure is not a rapid
interconversion of contributing structures. Several Lewis structures are
used together, because none of them exactly represents the actual
structure. To represent the intermediate, a resonance hybrid is used
instead.
- The
contributing structures are not isomers. They differ only in the position
of electrons, not in the position of nuclei. When transforming from one
Lewis structure into another, eventually no sigma bonds are broken (except
in the case of hyperconjugation). Only pi bonds differ.
- Each
Lewis formula must have the same number of valence electrons (and thus the
same total charge), and the same number of unpaired electrons, if any. [6]
- Bonds
that have different bond orders in different contributing structures do
not have typical bond lengths. Measurements reveal intermediate bond
lengths.
- The
real structure has a lower total potential energy than each of the
contributing structures would have. This means that it is more stable than
each separate contributing structure would be.
Resonance Energy:
Each structure is associated with a certain quantity of
energy, which determines the stability of the molecule or ion (the lower
energy, the greater stability). A resonance hybrid has a structure that is
intermediate between the contributing structures; the total quantity of
potential energy, however, is lower than the intermediate. Hybrids are
therefore always more stable than any of the contributing structures would be.[
The molecule is sometimes said to be "stabilized by resonance" or
"resonance-stabilized," but the stabilization derives from electron
delocalization, of which "resonance" is only a description.
Delocalization of the π-electrons lowers the orbital energies, imparting this
stability. The difference between the potential energy of the actual structure
(the resonance hybrid) and that of the contributing structure with the lowest
potential energy is called the "resonance energy".
Resonance energy of benzene:
Resonance (or delocalization) energy is the amount of energy
needed to convert the true delocalized structure into that of the most stable
contributing structure. The empirical resonance energy can be estimated by
comparing the heat of hydrogenation of the real substance with that estimated
for the contributing structure.
The complete hydrogenation of benzene to cyclohexane via
1,3-cyclohexadiene and cyclohexene is exothermic; 1 mole benzene delivers 208.4
kJ (49.8 kcal)
