The main points of molecular orbital theory and the principle of linear combination of atomic orbitals

Valence bond theory, hybrid orbital theory, and valence layer electron pair mutual exclusion theory all belong to modern valence bond theory.
It is believed that electrons move in atomic orbitals and electrons are localized
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1.
The main points of molecular orbital theory

(1) Molecular orbitals are linear combinations of atomic orbitals of atoms in molecules, and the number of molecular orbitals is equal to the number of atomic orbitals participating in the combination
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Ψ _MO = c ₁ Ψ _H1 + c ₂ Ψ _H2Ψ* _MO = c ₁ Ψ _H1 -c ₂ Ψ _H2

In the formula, c ₁ and c ₂ are constants
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(2) Among the molecular orbitals formed by linear combination of atomic orbitals, those with higher energy than the original atomic orbitals are called antibonding orbitals (Ψ* _MO ), and those with lower energy than the original atomic orbitals are called bonding orbitals (Ψ _MO ), and the energy is equal to The original atomic orbitals are called non-bonded orbitals
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(3) Each molecular orbital has its own wave function and its own angle distribution diagram
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(4) All electrons in a molecule belong to the whole molecule, and the arrangement of electrons in molecular orbitals also follows the principle of lowest energy, Pauli's principle and Hund's rule
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ss combination: the linear combination of the 1s orbitals of two atoms forms the bonding molecular orbital σ _1s and the anti-bonding molecular orbital σ* _1s ; if the 2s atomic orbital is linearly combined, it is combined into the bonding molecular orbital σ _2s and the anti-bonding molecular orbital σ* _2s
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The linear combination of two atomic orbitals forms a bonded molecular orbital and an anti-bonded molecular orbital
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pp combination: There are two ways of linear combination of the p orbitals of two atoms, namely "head-to-head" and "side-by-side".
The former is combined into a σ orbital, and the latter is combined into a π orbital
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Figure 6-20 Molecular orbitals of ss combination

When two atoms overlap along the x-axis, the two p _x and orbitals form the "head-to-head" bonding molecular orbital σ _px and anti-bonding molecular orbital σ* _px (Figure 6-21); at the same time, the p _{y of the} two atoms When the orbitals or p _z orbitals overlap, the "side-by-side" bonding molecular orbitals π _py , or πpz and anti-bonding molecular orbitals π* _py or πp*z are formed respectively (Figure 6-22)
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Figure 6-21 Molecular orbitals of pp "head meet" combination

Figure 6-22 Molecular orbitals of pp "side-by-side" combination

From Figure 6-20, Figure 6-21, and Figure 6-22, we can see that there are no nodal planes between the two nuclei of the bonding molecular orbital, and there are nodal planes between the two nuclei of the anti-bonding molecular orbital; the π molecular orbital has nodal planes passing through the bond axis.

2.

Molecular orbitals are formed by linear combinations of atomic orbitals.

1) Symmetry matching principle

According to the relationship between overlapping orbitals and bond axes, only atomic orbitals with the same symmetry can be combined to form molecular orbitals
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In the valence bond theory, it was introduced that when approaching along the x-axis, the s orbital and p, orbital symmetry are the same to form a σ bond, and p _y and p _y , p _z and p _z form a π bond
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According to molecular orbital theory, when the x-axis is used as the bond axis to combine into sub-orbitals, ss, sp _x and p _x -p _{x are} combined into σ orbitals, p _y- p _y , p _z- p _z , p _y- d _xy , p _z -d _xz combine to form a π orbital
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2) The principle of similar energy

Atomic orbitals with similar energies can be combined into effective molecular orbitals, and the closer the energy of the atomic orbitals, the lower the energy of the combined molecular orbitals, and the more stable the chemical bond formed
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The energy of the 1s orbital of the H atom (-1312kJ·mol ^-1 ) is similar to the energy of the 2p orbital of the O atom (-1251kJ·mol ^-1 ) and the energy of the 3p orbital of the Cl atom (-1314kJ·mol ^-1 ).
H, O, Cl can be combined into effective molecular orbitals, such as the formation of HCl, H ₂ O, ClO ₂ and other covalent compounds; while the energy of the 3s orbital of Na atom (-496kJ·mol ^-1 ) is relatively high, and cannot be combined with H, O, Cl combines into effective molecular orbitals, and Na can only form ionic bonds with H, O, and Cl.
For example, NaCl, NaH, and NaO ₂ are all ionic compounds
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3) The principle of maximum overlap of orbits

On the basis of satisfying the principle of symmetry matching and similar energy, the greater the degree of atomic orbital overlap, the lower the energy of the combined molecular orbital, and the more stable the chemical bond formed
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If two atomic orbitals overlap into sub-orbitals along the x-axis, p _x -p _x overlaps as "head-to-head", and p _y -p _y and p _z -p _z overlap as "side by side".
Generally speaking, between p orbitals The degree of overlap of "head-to-head" is greater than that of "side by side", that is, the σ bond is more stable than the π bond
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Related link: The relationship between hybrid orbital theory and the theory of mutual exclusion of valence electron pairs