How To Draw Molecular Orbital Diagram Khan Academy

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 general and the linear combination of atomic orbitals (LCAO) method in particular.^{[1] [2] [iii]} A primal principle of these theories is that as atoms bond to form molecules, a certain number of atomic orbitals combine to form the aforementioned 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 than complex when discussing even comparatively simple polyatomic molecules, such every bit marsh gas. MO diagrams can explain why some molecules exist and others do not. They can also predict bond force, as well as the electronic transitions that tin can have place.

History [edit]

Qualitative MO theory was introduced in 1928 past Robert S. Mulliken^{[iv] [5]} and Friedrich Hund.^[6] A mathematical description was provided by contributions from Douglas Hartree in 1928^[7] 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 center, flanked by constituent atomic orbital (AO) free energy levels for comparison, with the energy levels increasing from the bottom to the height. Lines, ofttimes dashed diagonal lines, connect MO levels with their constituent AO levels. Degenerate free 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-scale vertical arrows whose directions signal 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 ii atoms, whose AO unbonded energies are shown on the sides. For uncomplicated polyatomic molecules with a "key atom" such as methyl hydride (CH
₄ ) or carbon dioxide (CO
_ii ), a MO diagram may show i of the identical bonds to the central cantlet. For other polyatomic molecules, an MO diagram may testify one or more 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 exist omitted from a diagram for simplicity.

In MO theory molecular orbitals form by the overlap of atomic orbitals. Because 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 free energy correlates with electronegativity as more than electronegative atoms agree their electrons more tightly, lowering their energies. Sharing of molecular orbitals between atoms is more important when the diminutive orbitals have comparable energy; when the energies differ greatly the orbitals tend to be localized on one atom and the mode of bonding becomes ionic. A 2nd status for overlapping atomic orbitals is that they have the aforementioned symmetry.

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

Two atomic orbitals can overlap in two ways depending on their stage relationship (or relative signs for real orbitals). The stage (or sign) of an orbital is a directly outcome of the wave-similar 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 past shading one lobe. The sign of the phase itself does not have concrete meaning except when mixing orbitals to form molecular orbitals.

Two same-sign orbitals accept a constructive overlap forming a molecular orbital with the majority of the electron density located between the 2 nuclei. This MO is called the bonding orbital and its free energy is lower than that of the original atomic orbitals. A bail involving molecular orbitals which are symmetric with respect to whatsoever rotation around the bond axis is chosen a sigma bond (σ-bond). If the stage cycles once while rotating circular the axis, the bond is a pi bond (π-bond). Symmetry labels are further defined by whether the orbital maintains its original character after an inversion nigh its center; if it does, it is defined gerade, g. If the orbital does not maintain its original grapheme, it is ungerade, u.

Atomic orbitals can also interact with each other out-of-phase which leads to subversive counterfoil and no electron density betwixt the two nuclei at the so-called nodal plane depicted as a perpendicular dashed line. In this anti-bonding MO with free energy much higher than the original AO'southward, any electrons present are located in lobes pointing abroad from the primal internuclear axis. For a corresponding σ-bonding orbital, such an orbital would be symmetrical but differentiated from information technology by an asterisk as in σ*. For a π-bond, respective bonding and antibonding orbitals would not have such symmetry around the bond centrality and be designated π and π*, respectively.

The next footstep in constructing 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 everyman free energy
• The Pauli exclusion principle states that the maximum number of electrons occupying an orbital is two, with opposite spins
• Hund'southward rule states that when there are several MO's with equal energy, the electrons occupy the MO'southward one at a time earlier two electrons occupy the same MO.

The filled MO highest in energy is called the highest occupied molecular orbital (Human) and the empty MO just 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 forcefulness for chemical bail formation. Whenever mixing for an atomic orbital is not possible for reasons of symmetry or energy, a non-bonding MO is created, which is often quite like to and has energy level equal or close to its constituent AO, thus not contributing to bonding energetics. The resulting electron configuration can be described in terms of bail type, parity and occupancy for instance dihydrogen 1σ _g ². Alternatively it can be written equally a molecular term symbol eastward.g. ⁱΣ_k ⁺ for dihydrogen. Sometimes, the letter n is used to designate a non-bonding orbital.

For a stable bond, the bail club divers every bit

${\displaystyle \ {\mbox{bail society}}={\frac {({\mbox{number of electrons in bonding MOs}})-({\mbox{number of electrons in antibonding MOs}})}{2}}}$

must exist positive.

The relative order in MO energies and occupancy corresponds with electronic transitions establish in photoelectron spectroscopy (Pes). In this style it is possible to experimentally verify MO theory. In general, sharp Pes transitions indicate nonbonding electrons and wide bands are indicative of bonding and antibonding delocalized electrons. Bands tin resolve into fine structure with spacings corresponding to vibrational modes of the molecular cation (see Franck–Condon principle). PES energies are dissimilar from ionisation energies which relates to the energy required to strip off the nth electron afterwards the showtime northward − 1 electrons take been removed. MO diagrams with energy values 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 aforementioned symmetry formed from the combination of 2s and 2p diminutive orbitals are close plenty in energy to farther interact, which tin atomic number 82 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 atomic orbitals. By and large, in order to predict their relative energies, it is sufficient to consider but one atomic 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_g MO tin be roughly considered to be formed from interaction of oxygen 2p_z AOs only. It is found to exist lower in energy than the 1π_u MO, both experimentally and from more than sophisticated computational models, so that the expected order of filling is the 3σ_g before the 1π_u.^[ten] Hence the approximation to ignore the effects of further interactions is valid. Yet, experimental and computational results for homonuclear diatomics from Li₂ to N_two and certain heteronuclear combinations such as CO and NO testify that the 3σ_yard MO is higher in free energy than (and therefore filled afterward) the 1π_u MO.^[11] This tin be rationalised as the beginning-approximation 3σ_g has a suitable symmetry to interact with the 2σ_thou bonding MO formed from the 2s AOs. Every bit a result, the 2σ_yard is lowered in energy, whilst the 3σ_g is raised. For the aforementioned molecules this results in the 3σ_thou being higher in energy than the 1π_u MO, which is where s-p mixing is most evident. Too, interaction between the 2σ_u* and 3σ_u* MOs leads to a lowering in energy of the quondam and a raising in energy of the latter.^[9] However this is of less significance than the interaction of the bonding MOs.

Diatomic MO diagrams [edit]

A diatomic molecular orbital diagram is used to empathize the bonding of a diatomic molecule. MO diagrams tin can exist used to deduce magnetic backdrop of a molecule and how they alter with ionization. They likewise give insight to the bond order of the molecule, how many bonds are shared betwixt the two atoms.^[12]

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

Diatomic molecules consist of a bail between just two atoms. They can be cleaved into two categories: homonuclear and heteronuclear. A homonuclear diatomic molecule is one equanimous of two atoms of the same element. Examples are H₂, O_two, and N₂. A heteronuclear diatomic molecule is composed of two atoms of two different elements. Examples include CO, HCl, and NO.

Dihydrogen [edit]

Bail breaking in MO diagram

The smallest molecule, hydrogen gas exists equally dihydrogen (H-H) with a unmarried covalent bond between 2 hydrogen atoms. As each hydrogen atom has a single 1s atomic orbital for its electron, the bail forms past overlap of these 2 atomic orbitals. In the figure the two diminutive orbitals are depicted on the left and on the right. The vertical centrality always represents the orbital energies. Each atomic orbital is singly occupied with an up or downward arrow representing an electron.

Awarding of MO theory for dihydrogen results in having both electrons in the bonding MO with electron configuration 1σ _m ^two. The bail order for dihydrogen is (two-0)/2 = 1. The photoelectron spectrum of dihydrogen shows a single gear up of multiplets between 16 and 18 eV (electron volts).^[14]

The dihydrogen MO diagram helps explain how a bond breaks. When applying energy to dihydrogen, a molecular electronic transition takes identify when one electron in the bonding MO is promoted to the antibonding MO. The outcome is that at that place is no longer a net proceeds in energy.

The superposition of the two 1s atomic orbitals leads to the formation of the σ and σ* molecular orbitals. Two atomic orbitals in stage create a larger electron density, which leads to the σ orbital. If the two 1s orbitals are not in stage, a node betwixt them causes a jump in energy, the σ* orbital. From the diagram you can deduce the bond order, how many bonds are formed between the two atoms. For this molecule it is equal to one. Bond order can also give insight to how close or stretched a bond 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, but each helium has two electrons in its 1s diminutive orbital rather than one for hydrogen, then there are now iv electrons to place in the newly formed molecular orbitals.

The only fashion to accomplish this is by occupying both the bonding and antibonding orbitals with two electrons, which reduces the bond order ((2−2)/2) to goose egg and cancels the net free energy stabilization. Notwithstanding, by removing one electron from dihelium, the stable gas-stage species He ⁺
_two ion is formed with bond order ane/2.

Another molecule that is precluded based on this principle is diberyllium. Glucinium has an electron configuration 1s²2sⁱⁱ, and then there are again two electrons in the valence level. Nevertheless, the 2s can mix with the 2p orbitals in diberyllium, whereas there 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 net consequence 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.^{[15] [16]} The slight bonding nature explains the depression dissociation energy of merely 59 kJ·mol^−one.^[fifteen]

Dilithium [edit]

MO theory correctly predicts that dilithium is a stable molecule with bail order 1 (configuration 1σ _g ^two1σ _u ⁱⁱ2σ _yard ²). The 1s MOs are completely filled and do not participate in bonding.

Dilithium is a gas-phase molecule with a much lower bail strength than dihydrogen considering the 2s electrons are further removed from the nucleus. In a more detailed analysis^[17] which considers the environment 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 (encounter table 1).

Diboron [edit]

The MO diagram for diboron (B-B, electron configuration 1σ _g ²1σ _u ²2σ _m ^two2σ _u ²1π _u ²) requires the introduction of an atomic orbital overlap model for p orbitals. The three dumbbell-shaped p-orbitals accept 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's, the σ 2p has some not-bonding electron density at either side of the nuclei and the σ* 2p has some electron density between 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 two lobes above and beneath the airplane of the molecule. The orbital is not symmetric around the molecular axis and is therefore a pi orbital. The antibonding pi orbital (also asymmetrical) has four lobes pointing abroad from the nuclei. Both p_y and p_ten orbitals form a pair of pi orbitals equal in free energy (degenerate) and can have higher or lower energies than that of the sigma orbital.

In diboron the 1s and 2s electrons do not participate in bonding but the single electrons in the 2p orbitals occupy the 2πp_y and the 2πp_x MO'south resulting in bond gild 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 ²2σ_thousand ²2σ_u ^two1π_u ⁴) is a reactive gas-phase molecule. The molecule can exist described as having two pi bonds simply without a sigma bond.^[18]

Dinitrogen [edit]

Molecular orbital diagram of dinitrogen

With nitrogen, we run into 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, aforementioned with the 2s σ. This likewise causes a large bound in free energy in the 2p σ* orbital. The bail order of diatomic nitrogen is three, and information technology is a diamagnetic molecule.^[12]

The bail order for dinitrogen (1σ_g ²1σ_u ²2σ_g ²2σ_u ^two1π_u ⁴3σ_g ²) is three because two electrons are at present too 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 summit 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_{one thousand} ^two at fifteen.5 eV (sharp).

Dioxygen [edit]

Molecular orbital diagram of dioxygen

Oxygen has a similar setup to H₂, but now we consider 2s and 2p orbitals. When creating the molecular orbitals from the p orbitals, notice the three diminutive orbitals dissever into 3 molecular orbitals, a singly degenerate σ and a doubly degenerate π orbital. Another property we can observe by examining molecular orbital diagrams is the magnetic property of diamagnetic or paramagnetic. If all the electrons are paired, at that place 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 example of a paramagnetic diatomic. As well discover the bond order of diatomic oxygen is two. ^[12]

MO treatment of dioxygen is dissimilar from that of the previous diatomic molecules considering the pσ MO is now lower in energy than the 2π orbitals. This is attributed to interaction between the 2s MO and the 2p_z MO.^[20] Distributing 8 electrons over half-dozen molecular orbitals leaves the final two electrons as a degenerate pair in the 2pπ* antibonding orbitals resulting in a bond guild of 2. Equally in diboron, these two unpaired electrons have the same spin in the basis state, which is a paramagnetic diradical triplet oxygen. The first excited state has both Man electrons paired in one orbital with reverse spins, and is known as singlet oxygen.

MO diagram of dioxygen triplet basis state

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

Difluorine and dineon [edit]

In difluorine 2 additional electrons occupy the 2pπ* with a bail order of 1. In dineon Ne
₂ (every bit with dihelium) the number of bonding electrons equals the number of antibonding electrons and this molecule does non exist.

Dimolybdenum and ditungsten [edit]

MO diagram of dimolybdenum

Dimolybdenum (Mo₂) is notable for having a sextuple bond. This involves two sigma bonds (4d_zⁱⁱ and 5s), two pi bonds (using 4d_xz and 4d_yz), and 2 delta bonds (4d_{10² − y²} and 4d_xy). Ditungsten (West₂) has a similar structure.^{[21] [22]}

MO energies overview [edit]

Table 1 gives an overview of MO energies for beginning 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	C2	N2	O2	F2
1σg	-0.5969	-two.4523	-7.7040	- 11.3598	- 15.6820	- xx.7296	-26.4289
1σu		-2.4520	-vii.7032	-11.3575	-15.6783	-20.7286	-26.4286
2σone thousand		-0.1816	-0.7057	-1.0613	-1.4736	-1.6488	-1.7620
2σu			-0.3637	-0.5172	-0.7780	-1.0987	-i.4997
3σ1000					-0.6350	-0.7358	-0.7504
1πu			-0.3594	-0.4579	-0.6154	-0.7052	-0.8097
1πg						-0.5319	-0.6682
1s (AO)	-0.5	-2.4778	-7.6953	-11.3255	-fifteen.6289	-20.6686	-26.3829
2s (AO)		-0.1963	-0.4947	-0.7056	-0.9452	-i.2443	-1.5726
2p (AO)			-0.3099	-0.4333	-0.5677	-0.6319	-0.7300

Heteronuclear diatomics [edit]

In heteronuclear diatomic molecules, mixing of atomic orbitals merely 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 caste of mixing is low. The electron configuration 1σⁱⁱ1σ*ⁱⁱ2σⁱⁱ2σ*²1π⁴3σ² is identical to that of nitrogen. The one thousand and u subscripts no longer apply because the molecule lacks a center of symmetry.

In hydrogen fluoride (HF), the hydrogen 1s orbital tin can mix with fluorine 2p_z orbital to form a sigma bond because experimentally the energy of 1s of hydrogen is comparable with 2p of fluorine. The HF electron configuration 1σ^two2σ²3σⁱⁱ1π^iv reflects that the other electrons remain in 3 alone pairs and that the bond guild is 1.

The more than electronegative atom is the more energetically excited because information technology more similar in energy to its atomic orbital. This also accounts for the majority of the electron negativity residing effectually the more electronegative molecule. Applying the LCAO-MO method allows us to move away from a more than static Lewis structure type approach and really business relationship for periodic trends that influence electron motility. Non-bonding orbitals refer to lone pairs seen on certain atoms in a molecule. A farther understanding for the energy level refinement can be caused by delving into quantum chemistry; the Schrödinger equation can be applied to predict movement 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 aforementioned as for the homonuclear molecules. It has a bond order of 2.5 and is a paramagnetic molecule. The free energy differences of the 2s orbitals are different enough that each produces its ain non-bonding σ orbitals. Discover this is a good instance of making the ionized NO⁺ stabilize the bond and generate a triple bond, also changing the magnetic holding to diamagnetic.^[12]

HF [edit]

Molecular orbital diagram of HF

Hydrogen fluoride is another case of a heteronuclear molecule. Information technology is slightly different in that the π orbital is non-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 bond order of one.

Triatomic molecules [edit]

Carbon dioxide [edit]

Carbon dioxide, CO
₂ , is a linear molecule with a total of sixteen bonding electrons in its valence shell. Carbon is the fundamental atom of the molecule and a principal axis, the z-axis, is visualized as a unmarried centrality that goes through the center of carbon and the 2 oxygens atoms. For convention, blue 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 (−19.four eV), carbon 2p (−10.seven eV), and oxygen 2p (−xv.9 eV)) energies associated with the atomic orbitals are in proximity whereas the oxygen 2s energy (−32.4 eV) is unlike.^[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 atomic orbitals, symmetry labels are deduced with respect to rotation about the principal axis which generates a phase modify, pi bond (π)^[26] or generates no stage change, known as a sigma bond (σ).^[27] Symmetry labels are farther divers past whether the diminutive orbital maintains its original graphic symbol after an inversion about its center atom; if the atomic orbital does retain its original grapheme it is divers gerade, g, 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'due south molecular orbitals are made by the linear combination of atomic orbitals of the aforementioned irreducible representation that are also similar in atomic orbital energy. Significant atomic orbital overlap explains why sp bonding may occur.^[28] Strong mixing of the oxygen 2s atomic orbital is non to exist expected and are non-bonding degenerate molecular orbitals. The combination of like atomic orbital/wave functions and the combinations of atomic orbital/wave function inverses create particular energies associated with the nonbonding (no modify), bonding (lower than either parent orbital energy) and antibonding (higher free energy than either parent atomic 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 non σ or π but depend on the symmetry of each molecule. H2o (H
_two O) is a bent molecule (105°) with C_2v molecular symmetry. The possible orbital symmetries are listed in the table below. For instance, an orbital of B_i symmetry (called a b₁ orbital with a small-scale b since it is a i-electron part) is multiplied by -ane under the symmetry operations C₂ (rotation nigh the 2-fold rotation axis) and σ_v'(yz) (reflection in the molecular plane). Information technology is multiplied by +one(unchanged) by the identity operation E and by σ_v(xz) (reflection in the plane bisecting the H-O-H bending).

Molecular orbital diagram of h2o

C2v	East	Cii	σv(xz)	σ5'(yz)
A1	i	one	1	1	z	x two, y two, z 2
A2	i	1	−i	−one	Rz	xy
B1	i	−1	1	−ane	10, Ry	xz
B2	1	−1	−1	1	y, Rx	yz

The oxygen atomic orbitals are labeled co-ordinate to their symmetry as a_ane for the 2s orbital and b_one (2p_ten), b₂ (2p_y) and a_ane (2p_z) for the three 2p orbitals. The two hydrogen 1s orbitals are premixed to form a_ane (σ) and b_two (σ*) MO.

Mixing takes place between aforementioned-symmetry orbitals of comparable energy resulting a new gear up of MO'due south for h2o:

• 2a_ane MO from mixing of the oxygen 2s AO and the hydrogen σ MO.
• 1b_two 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 agreement with this description the photoelectron spectrum for water shows a sharp peak for the nonbonding 1b₁ MO (12.half-dozen eV) and three broad peaks for the 3a_one MO (fourteen.vii eV), 1b_ii MO (18.five eV) and the 2a₁ MO (32.two eV).^[29] The 1b₁ MO is a lone pair, while the 3a₁, 1b_ii and 2a_one MO's can exist localized to give two O−H bonds and an in-aeroplane solitary pair.^[thirty] This MO treatment of h2o does non accept ii equivalent rabbit ear alone pairs.^[31]

Hydrogen sulfide (H₂S) too has a C_2v symmetry with eight valence electrons but the bending bending is only 92°. As reflected in its photoelectron spectrum as compared to water the 5a₁ MO (corresponding to the 3a_one MO in water) is stabilised (improved overlap) and the 2b₂ MO (respective to the 1b₂ 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

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

Posted by: scotttheatione.blogspot.com

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