How to Read a Molecular Orbital Configuration

Visual tool in quantum chemistry

A molecular orbital diagram, or MO diagram, is a qualitative descriptive tool explaining chemical bonding in molecules in terms of molecular orbital theory in full general and the linear combination of atomic orbitals (LCAO) method in particular.^{[1] [two] [3]} A central principle of these theories is that as atoms bond to form molecules, a sure number of atomic orbitals combine to form the same number of molecular orbitals, although the electrons involved may be redistributed amidst the orbitals. This tool is very well suited for simple diatomic molecules such as dihydrogen, dioxygen, and carbon monoxide but becomes more complex when discussing even insufficiently simple polyatomic molecules, such equally methane. MO diagrams can explain why some molecules exist and others practice not. They can also predict bond strength, as well as the electronic transitions that can have place.

History [edit]

Qualitative MO theory was introduced in 1928 by Robert South. Mulliken^{[4] [five]} and Friedrich Hund.^[6] A mathematical description was provided by contributions from Douglas Hartree in 1928^[seven] and Vladimir Fock in 1930.^[8]

Basics [edit]

Molecular orbital diagrams are diagrams of molecular orbital (MO) energy levels, shown as short horizontal lines in the eye, flanked by elective diminutive orbital (AO) energy levels for comparison, with the energy levels increasing from the lesser to the top. Lines, oft dashed diagonal lines, connect MO levels with their constituent AO levels. Degenerate energy levels are commonly shown side by side. Appropriate AO and MO levels are filled with electrons past the Pauli Exclusion Principle, symbolized by small vertical arrows whose directions indicate the electron spins. The AO or MO shapes themselves are often non shown on these diagrams. For a diatomic molecule, an MO diagram effectively shows the energetics of the bond between the two atoms, whose AO unbonded energies are shown on the sides. For simple polyatomic molecules with a "central atom" such as methyl hydride (CH
₄ ) or carbon dioxide (CO
_two ), a MO diagram may show one of the identical bonds to the primal atom. For other polyatomic molecules, an MO diagram may show one or more than bonds of interest in the molecules, leaving others out for simplicity. Often even for simple molecules, AO and MO levels of inner orbitals and their electrons may be omitted from a diagram for simplicity.

In MO theory molecular orbitals class by the overlap of atomic orbitals. Considering σ bonds characteristic greater overlap than π bonds, σ bonding and σ* antibonding orbitals feature greater energy splitting (separation) than π and π* orbitals. The atomic orbital energy correlates with electronegativity as more electronegative atoms concord their electrons more tightly, lowering their energies. Sharing of molecular orbitals between atoms is more important when the atomic orbitals have comparable free energy; when the energies differ profoundly the orbitals tend to be localized on i atom and the way of bonding becomes ionic. A second status for overlapping atomic orbitals is that they have the aforementioned symmetry.

MO diagram for dihydrogen. Here electrons are shown past dots.

Two atomic orbitals can overlap in 2 ways depending on their stage relationship (or relative signs for real orbitals). The phase (or sign) of an orbital is a directly issue of the wave-like properties of electrons. In graphical representations of orbitals, orbital sign is depicted either by a plus or minus sign (which has no relationship to electric charge) or by shading one lobe. The sign of the phase itself does non have concrete significant except when mixing orbitals to form molecular orbitals.

Two same-sign orbitals have a constructive overlap forming a molecular orbital with the bulk of the electron density located betwixt the two nuclei. This MO is called the bonding orbital and its free energy is lower than that of the original atomic orbitals. A bond involving molecular orbitals which are symmetric with respect to whatsoever rotation around the bail centrality is called a sigma bail (σ-bond). If the phase cycles one time while rotating round the axis, the bond is a pi bond (π-bail). Symmetry labels are farther defined by whether the orbital maintains its original character later on an inversion about its center; if it does, it is defined gerade, thou. If the orbital does not maintain its original grapheme, it is ungerade, u.

Diminutive orbitals can also interact with each other out-of-phase which leads to destructive cancellation and no electron density between the 2 nuclei at the and so-called nodal plane depicted as a perpendicular dashed line. In this anti-bonding MO with energy much higher than the original AO's, any electrons present are located in lobes pointing abroad from the central internuclear axis. For a corresponding σ-bonding orbital, such an orbital would exist symmetrical merely differentiated from it by an asterisk as in σ*. For a π-bond, corresponding bonding and antibonding orbitals would not take such symmetry around the bail axis and be designated π and π*, respectively.

The side by side step in amalgam an MO diagram is filling the newly formed molecular orbitals with electrons. Three general rules utilise:

• The Aufbau principle states that orbitals are filled starting with the lowest energy
• The Pauli exclusion principle states that the maximum number of electrons occupying an orbital is 2, with reverse spins
• Hund'south rule states that when there are several MO's with equal energy, the electrons occupy the MO's one at a time earlier two electrons occupy the same MO.

The filled MO highest in energy is chosen the highest occupied molecular orbital (Human) and the empty MO only above it is then the lowest unoccupied molecular orbital (LUMO). The electrons in the bonding MO's are called bonding electrons and whatever electrons in the antibonding orbital would be called antibonding electrons. The reduction in energy of these electrons is the driving strength for chemical bond formation. Whenever mixing for an diminutive orbital is not possible for reasons of symmetry or energy, a non-bonding MO is created, which is frequently quite similar to and has energy level equal or close to its constituent AO, thus not contributing to bonding energetics. The resulting electron configuration tin can be described in terms of bond type, parity and occupancy for example dihydrogen 1σ _g ⁱⁱ. Alternatively it can be written every bit a molecular term symbol e.g. ¹Σ_k ⁺ for dihydrogen. Sometimes, the letter n is used to designate a not-bonding orbital.

For a stable bond, the bond order divers as

${\displaystyle \ {\mbox{bond order}}={\frac {({\mbox{number of electrons in bonding MOs}})-({\mbox{number of electrons in antibonding MOs}})}{ii}}}$

must exist positive.

The relative guild in MO energies and occupancy corresponds with electronic transitions establish in photoelectron spectroscopy (PES). In this way information technology is possible to experimentally verify MO theory. In full general, sharp PES transitions indicate nonbonding electrons and broad bands are indicative of bonding and antibonding delocalized electrons. Bands can resolve into fine structure with spacings corresponding to vibrational modes of the molecular cation (see Franck–Condon principle). Pes energies are different from ionisation energies which relates to the energy required to strip off the nth electron subsequently the first n − 1 electrons take been removed. MO diagrams with energy values tin can be obtained mathematically using the Hartree–Fock method. The starting point for any MO diagram is a predefined molecular geometry for the molecule in question. An exact relationship between geometry and orbital energies is given in Walsh diagrams.

s-p mixing [edit]

The phenomenon of s-p mixing occurs when molecular orbitals of the same symmetry formed from the combination of 2s and 2p atomic orbitals are close enough in energy to further collaborate, which tin can lead to a change in the expected order of orbital energies.^[9] When molecular orbitals are formed, they are mathematically obtained from linear combinations of the starting diminutive orbitals. Generally, in order to predict their relative energies, it is sufficient to consider simply 1 diminutive orbital from each atom to grade a pair of molecular orbitals, as the contributions from the others are negligible. For instance, in dioxygen the 3σ_yard MO can exist roughly considered to be formed from interaction of oxygen 2p_z AOs only. It is establish to be lower in free energy than the 1π_u MO, both experimentally and from more sophisticated computational models, then that the expected order of filling is the 3σ_g before the 1π_u.^[10] Hence the approximation to ignore the effects of further interactions is valid. However, experimental and computational results for homonuclear diatomics from Li₂ to N₂ and certain heteronuclear combinations such as CO and NO show that the 3σ_g MO is higher in energy than (and therefore filled after) the 1π_u MO.^[11] This can exist rationalised as the showtime-approximation 3σ_g has a suitable symmetry to interact with the 2σ_g bonding MO formed from the 2s AOs. As a event, the 2σ_{one thousand} is lowered in energy, whilst the 3σ_g is raised. For the aforementioned molecules this results in the 3σ_g beingness higher in energy than the 1π_u MO, which is where due south-p mixing is well-nigh evident. As well, interaction between the 2σ_u* and 3σ_u* MOs leads to a lowering in free energy of the former and a raising in energy of the latter.^[9] Still this is of less significance than the interaction of the bonding MOs.

Diatomic MO diagrams [edit]

A diatomic molecular orbital diagram is used to understand the bonding of a diatomic molecule. MO diagrams can be used to deduce magnetic backdrop of a molecule and how they change with ionization. They also requite insight to the bail order of the molecule, how many bonds are shared between the ii atoms.^[12]

The energies of the electrons are further understood past applying the Schrödinger equation to a molecule. Quantum Mechanics is able to describe the energies exactly for unmarried electron systems merely tin can be approximated precisely for multiple electron systems using the Born-Oppenheimer Approximation, such that the nuclei are assumed stationary. The LCAO-MO method is used in conjunction to further describe the state of the molecule. ^[xiii]

Diatomic molecules consist of a bond between only two atoms. They can be broken into two categories: homonuclear and heteronuclear. A homonuclear diatomic molecule is one composed of two atoms of the same chemical element. Examples are H_ii, O₂, and North₂. A heteronuclear diatomic molecule is composed of 2 atoms of two different elements. Examples include CO, HCl, and NO.

Dihydrogen [edit]

Bond breaking in MO diagram

The smallest molecule, hydrogen gas exists as dihydrogen (H-H) with a single covalent bond betwixt two hydrogen atoms. Every bit each hydrogen atom has a single 1s atomic orbital for its electron, the bond forms by overlap of these two atomic orbitals. In the effigy the two atomic orbitals are depicted on the left and on the correct. The vertical axis always represents the orbital energies. Each atomic orbital is singly occupied with an upward or down pointer representing an electron.

Awarding of MO theory for dihydrogen results in having both electrons in the bonding MO with electron configuration 1σ ₁₀₀₀ ². The bond order for dihydrogen is (2-0)/2 = 1. The photoelectron spectrum of dihydrogen shows a single fix of multiplets betwixt 16 and 18 eV (electron volts).^[xiv]

The dihydrogen MO diagram helps explain how a bail breaks. When applying energy to dihydrogen, a molecular electronic transition takes place when one electron in the bonding MO is promoted to the antibonding MO. The result is that there is no longer a internet gain in energy.

The superposition of the two 1s diminutive orbitals leads to the formation of the σ and σ* molecular orbitals. Two atomic orbitals in phase create a larger electron density, which leads to the σ orbital. If the two 1s orbitals are not in stage, a node between them causes a spring in energy, the σ* orbital. From the diagram you can deduce the bond order, how many bonds are formed betwixt the ii atoms. For this molecule it is equal to i. Bond gild can too give insight to how close or stretched a bail has become if a molecule is ionized.^[12]

Dihelium and diberyllium [edit]

Dihelium (He-He) is a hypothetical molecule and MO theory helps to explain why dihelium does not exist in nature. The MO diagram for dihelium looks very similar to that of dihydrogen, only each helium has 2 electrons in its 1s atomic orbital rather than one for hydrogen, and so there are now four electrons to place in the newly formed molecular orbitals.

The simply way to accomplish this is by occupying both the bonding and antibonding orbitals with 2 electrons, which reduces the bail order ((2−2)/2) to null and cancels the net energy stabilization. Still, by removing 1 electron from dihelium, the stable gas-stage species He ⁺
₂ ion is formed with bail order i/2.

Another molecule that is precluded based on this principle is diberyllium. Beryllium has an electron configuration 1s²2s², so at that place are once more two electrons in the valence level. Nonetheless, the 2s can mix with the 2p orbitals in diberyllium, whereas in that location are no p orbitals in the valence level of hydrogen or helium. This mixing makes the antibonding 1σ_u orbital slightly less antibonding than the bonding 1σ_g orbital is bonding, with a internet event that the whole configuration has a slight bonding nature. This explains the fact that the diberyllium molecule exists and has been observed in the gas phase.^{[xv] [xvi]} The slight bonding nature explains the low dissociation free energy of only 59 kJ·mol⁻¹.^[15]

Dilithium [edit]

MO theory correctly predicts that dilithium is a stable molecule with bond society 1 (configuration 1σ _thou ²1σ _u ²2σ _g ⁱⁱ). The 1s MOs are completely filled and do non participate in bonding.

Dilithium is a gas-phase molecule with a much lower bond strength than dihydrogen because the 2s electrons are farther removed from the nucleus. In a more detailed analysis^[17] which considers the surround of each orbital due to all other electrons, both the 1σ orbitals have higher energies than the 1s AO and the occupied 2σ is also higher in energy than the 2s AO (run across tabular array one).

Diboron [edit]

The MO diagram for diboron (B-B, electron configuration 1σ _chiliad ^two1σ _u ⁱⁱ2σ _g ^two2σ _u ²1π _u ²) requires the introduction of an atomic orbital overlap model for p orbitals. The 3 dumbbell-shaped p-orbitals have equal energy and are oriented mutually perpendicularly (or orthogonally). The p-orbitals oriented in the z-direction (p_z) can overlap end-on forming a bonding (symmetrical) σ orbital and an antibonding σ* molecular orbital. In contrast to the sigma 1s MO'southward, the σ 2p has some non-bonding electron density at either side of the nuclei and the σ* 2p has some electron density betwixt the nuclei.

The other two p-orbitals, p_y and p_x, can overlap side-on. The resulting bonding orbital has its electron density in the shape of ii lobes higher up and below the aeroplane of the molecule. The orbital is non symmetric effectually the molecular axis and is therefore a pi orbital. The antibonding pi orbital (also asymmetrical) has iv lobes pointing away from the nuclei. Both p_y and p_x orbitals form a pair of pi orbitals equal in energy (degenerate) and can have higher or lower energies than that of the sigma orbital.

In diboron the 1s and 2s electrons do non participate in bonding but the single electrons in the 2p orbitals occupy the 2πp_y and the 2πp_x MO's resulting in bail order 1. Because the electrons have equal energy (they are degenerate) diboron is a diradical and since the spins are parallel the molecule is paramagnetic.

In certain diborynes the boron atoms are excited and the bond order is 3.

Dicarbon [edit]

Like diboron, dicarbon (C-C electron configuration:1σ_g ²1σ_u ^two2σ₁₀₀₀ ²2σ_u ²1π_u ⁴) is a reactive gas-phase molecule. The molecule tin exist described as having 2 pi bonds but without a sigma bond.^[18]

Dinitrogen [edit]

Molecular orbital diagram of dinitrogen

With nitrogen, nosotros encounter the two molecular orbitals mixing and the energy repulsion. This is the reasoning for the rearrangement from a more familiar diagram. Notice how the σ from the 2p behaves more non-bonding like due to mixing, same with the 2s σ. This also causes a large jump in free energy in the 2p σ* orbital. The bond order of diatomic nitrogen is three, and information technology is a diamagnetic molecule.^[12]

The bond guild for dinitrogen (1σ₁₀₀₀ ^two1σ_u ²2σ_g ^two2σ_u ²1π_u ^four3σ_g ²) is three because ii electrons are now likewise added in the 3σ MO. The MO diagram correlates with the experimental photoelectron spectrum for nitrogen.^[nineteen] The 1σ electrons can be matched to a acme at 410 eV (broad), the 2σ_g electrons at 37 eV (broad), the 2σ_u electrons at xix eV (doublet), the 1π_u ⁴ electrons at 17 eV (multiplets), and finally the 3σ_g ² at 15.5 eV (sharp).

Dioxygen [edit]

Molecular orbital diagram of dioxygen

Oxygen has a similar setup to H_two, but at present we consider 2s and 2p orbitals. When creating the molecular orbitals from the p orbitals, detect the 3 diminutive orbitals split into three molecular orbitals, a singly degenerate σ and a doubly degenerate π orbital. Another belongings nosotros can notice by examining molecular orbital diagrams is the magnetic belongings of diamagnetic or paramagnetic. If all the electrons are paired, there is a slight repulsion and it is classified as diamagnetic. If unpaired electrons are present, it is attracted to a magnetic field, and therefore paramagnetic. Oxygen is an instance of a paramagnetic diatomic. Also notice the bail order of diatomic oxygen is ii. ^[12]

MO treatment of dioxygen is dissimilar from that of the previous diatomic molecules because the pσ MO is at present lower in energy than the 2π orbitals. This is attributed to interaction between the 2s MO and the 2p_z MO.^[xx] Distributing 8 electrons over 6 molecular orbitals leaves the final two electrons as a degenerate pair in the 2pπ* antibonding orbitals resulting in a bond order of 2. As in diboron, these ii unpaired electrons take the aforementioned spin in the ground country, which is a paramagnetic diradical triplet oxygen. The first excited state has both HOMO electrons paired in one orbital with opposite spins, and is known as singlet oxygen.

MO diagram of dioxygen triplet ground state

The bail order decreases and the bond length increases in the order O ⁺
₂ (112.two pm), O
₂ (121 pm), O ⁻
₂ (128 pm) and O ²⁻
₂ (149 pm).^[20]

Difluorine and dineon [edit]

In difluorine two additional electrons occupy the 2pπ* with a bond order of i. In dineon Ne
_two (as with dihelium) the number of bonding electrons equals the number of antibonding electrons and this molecule does not exist.

Dimolybdenum and ditungsten [edit]

MO diagram of dimolybdenum

Dimolybdenum (Mo_ii) is notable for having a sextuple bail. This involves two sigma bonds (4d_z² and 5s), two pi bonds (using 4d_xz and 4d_yz), and two delta bonds (4d_{x² − y²} and 4d_xy). Ditungsten (Westward_ii) has a similar structure.^{[21] [22]}

MO energies overview [edit]

Tabular array one gives an overview of MO energies for showtime row diatomic molecules calculated by the Hartree-Fock-Roothaan method, together with atomic orbital energies.

Table 1. Calculated MO energies for diatomic molecules in Hartrees [17]
	H2	Li2	B2	Cii	N2	O2	F2
1σthou	-0.5969	-ii.4523	-vii.7040	- 11.3598	- xv.6820	- 20.7296	-26.4289
1σu		-2.4520	-7.7032	-11.3575	-15.6783	-20.7286	-26.4286
2σg		-0.1816	-0.7057	-ane.0613	-one.4736	-1.6488	-i.7620
2σu			-0.3637	-0.5172	-0.7780	-ane.0987	-1.4997
3σm					-0.6350	-0.7358	-0.7504
1πu			-0.3594	-0.4579	-0.6154	-0.7052	-0.8097
1πthou						-0.5319	-0.6682
1s (AO)	-0.5	-2.4778	-seven.6953	-11.3255	-15.6289	-20.6686	-26.3829
2s (AO)		-0.1963	-0.4947	-0.7056	-0.9452	-i.2443	-one.5726
2p (AO)			-0.3099	-0.4333	-0.5677	-0.6319	-0.7300

Heteronuclear diatomics [edit]

In heteronuclear diatomic molecules, mixing of atomic orbitals only occurs when the electronegativity values are similar. In carbon monoxide (CO, isoelectronic with dinitrogen) the oxygen 2s orbital is much lower in energy than the carbon 2s orbital and therefore the degree of mixing is low. The electron configuration 1σ²1σ*²2σ²2σ*ⁱⁱ1π⁴3σ² is identical to that of nitrogen. The g and u subscripts no longer utilise because the molecule lacks a center of symmetry.

In hydrogen fluoride (HF), the hydrogen 1s orbital can mix with fluorine 2p_z orbital to course a sigma bond because experimentally the free energy of 1s of hydrogen is comparable with 2p of fluorine. The HF electron configuration 1σ²2σⁱⁱ3σⁱⁱ1π⁴ reflects that the other electrons remain in iii lone pairs and that the bond order is 1.

The more electronegative atom is the more energetically excited considering it more than similar in energy to its atomic orbital. This also accounts for the majority of the electron negativity residing around the more electronegative molecule. Applying the LCAO-MO method allows usa to move away from a more static Lewis structure type approach and actually account for periodic trends that influence electron move. Non-bonding orbitals refer to solitary pairs seen on certain atoms in a molecule. A further understanding for the free energy level refinement tin exist caused by delving into quantum chemistry; the Schrödinger equation can exist applied to predict move and describe the state of the electrons in a molecule.^{[13] [23]}

NO [edit]

Molecular orbital diagram of NO

Nitric oxide is a heteronuclear molecule that exhibits mixing. The construction of its MO diagram is the same as for the homonuclear molecules. It has a bond order of ii.v and is a paramagnetic molecule. The energy differences of the 2s orbitals are different enough that each produces its own not-bonding σ orbitals. Notice this is a adept example of making the ionized NO⁺ stabilize the bond and generate a triple bond, also irresolute the magnetic belongings to diamagnetic.^[12]

HF [edit]

Molecular orbital diagram of HF

Hydrogen fluoride is another example of a heteronuclear molecule. It is slightly different in that the π orbital is not-bonding, as well as the 2s σ. From the hydrogen, its valence 1s electron interacts with the 2p electrons of fluorine. This molecule is diamagnetic and has a bail order of one.

Triatomic molecules [edit]

Carbon dioxide [edit]

Carbon dioxide, CO
_ii , is a linear molecule with a full of sixteen bonding electrons in its valence beat out. Carbon is the fundamental cantlet of the molecule and a principal axis, the z-axis, is visualized equally a single axis that goes through the center of carbon and the 2 oxygens atoms. For convention, bluish atomic orbital lobes are positive phases, red diminutive orbitals are negative phases, with respect to the wave function from the solution of the Schrödinger equation.^[24] In carbon dioxide the carbon 2s (−nineteen.4 eV), carbon 2p (−10.vii eV), and oxygen 2p (−15.9 eV)) energies associated with the diminutive orbitals are in proximity whereas the oxygen 2s energy (−32.4 eV) is dissimilar.^[25]

Carbon and each oxygen atom volition have a 2s atomic orbital and a 2p atomic orbital, where the p orbital is divided into p_x, p_y, and p_z. With these derived diminutive orbitals, symmetry labels are deduced with respect to rotation about the principal axis which generates a phase alter, pi bond (π)^[26] or generates no phase change, known as a sigma bond (σ).^[27] Symmetry labels are further defined by whether the atomic orbital maintains its original character afterwards an inversion nigh its center atom; if the atomic orbital does retain its original character it is defined gerade, yard, or if the atomic orbital does not maintain its original character, ungerade, u. The final symmetry-labeled atomic orbital is now known as an irreducible representation.

Carbon dioxide's molecular orbitals are made past the linear combination of atomic orbitals of the same irreducible representation that are also like in atomic orbital energy. Significant diminutive orbital overlap explains why sp bonding may occur.^[28] Potent mixing of the oxygen 2s atomic orbital is non to exist expected and are non-bonding degenerate molecular orbitals. The combination of similar atomic orbital/wave functions and the combinations of atomic orbital/wave function inverses create detail energies associated with the nonbonding (no change), bonding (lower than either parent orbital energy) and antibonding (higher free energy than either parent diminutive orbital energy) molecular orbitals.

• MO model carbon dioxide

• 

  Atomic orbitals of carbon dioxide

• 

  Molecular orbitals of carbon dioxide

• 

  MO diagram of carbon dioxide

Water [edit]

For nonlinear molecules, the orbital symmetries are not σ or π but depend on the symmetry of each molecule. H2o (H
₂ O) is a aptitude molecule (105°) with C_2v molecular symmetry. The possible orbital symmetries are listed in the tabular array below. For example, an orbital of B₁ symmetry (called a b₁ orbital with a minor b since it is a one-electron function) is multiplied by -1 under the symmetry operations C_ii (rotation about the ii-fold rotation axis) and σ_v'(yz) (reflection in the molecular plane). It is multiplied by +ane(unchanged) by the identity operation E and by σ_five(xz) (reflection in the aeroplane bisecting the H-O-H bending).

Molecular orbital diagram of h2o

C2v	E	C2	σv(xz)	σ5'(yz)
A1	1	one	1	1	z	x two, y 2, z 2
A2	ane	ane	−i	−one	Rz	xy
Bane	1	−1	1	−ane	x, Ry	xz
B2	1	−1	−one	1	y, R10	yz

The oxygen atomic orbitals are labeled according to their symmetry as a_one for the 2s orbital and b_i (2p_x), b_two (2p_y) and a₁ (2p_z) for the three 2p orbitals. The two hydrogen 1s orbitals are premixed to form a₁ (σ) and b_two (σ*) MO.

Mixing takes place betwixt same-symmetry orbitals of comparable energy resulting a new gear up of MO's for water:

• 2a₁ MO from mixing of the oxygen 2s AO and the hydrogen σ MO.
• 1b₂ MO from mixing of the oxygen 2p_y AO and the hydrogen σ* MO.
• 3a₁ MO from mixing of the a₁ AOs.
• 1b₁ nonbonding MO from the oxygen 2p_x AO (the p-orbital perpendicular to the molecular plane).

In understanding with this description the photoelectron spectrum for h2o shows a precipitous top for the nonbonding 1b_i MO (12.6 eV) and iii broad peaks for the 3a₁ MO (14.vii eV), 1b₂ MO (18.5 eV) and the 2a₁ MO (32.2 eV).^[29] The 1b_one MO is a lone pair, while the 3a_i, 1b_ii and 2a₁ MO'southward tin can be localized to give two O−H bonds and an in-plane lone pair.^[30] This MO handling of h2o does non have 2 equivalent rabbit ear alone pairs.^[31]

Hydrogen sulfide (H₂S) likewise has a C_2v symmetry with eight valence electrons but the bending angle is just 92°. As reflected in its photoelectron spectrum as compared to water the 5a_i MO (corresponding to the 3a₁ MO in water) is stabilised (improved overlap) and the 2b₂ MO (respective to the 1b_two MO in h2o) is destabilized (poorer overlap).

References [edit]

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External links [edit]

• MO diagrams at meta-synthesis.com Link
• MO diagrams at chem1.com Link
• Molecular orbitals at winter.grouping.shef.ac.uk Link

johnsonhistogives.blogspot.com

Source: https://en.wikipedia.org/wiki/Molecular_orbital_diagram

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