Question: Who invented chemical bonds and molecular orbitals?
About the article "Chemistry", section "Chemical structure".

Candidate passages:
 "Not to be confused with s-p mixing in Molecular Orbital theory.
See Molecular orbital diagram.
In chemistry, hybridisation (or hybridization) is the concept of mixing atomic orbitals into new "hybrid orbitals" (with different energies, shapes, etc., than the component atomic orbitals) suitable for the pairing of electrons to form chemical bonds in valence bond theory.
Hybrid orbitals are very useful in the explanation of molecular geometry and atomic bonding properties.
Although sometimes taught together with the valence shell electron-pair repulsion (VSEPR) theory, valence bond and hybridisation are in fact not related to the VSEPR model. 
Bent's rule describes and explains the relationship between the orbital hybridisations of central atoms in molecules and the electronegativities of substituents.
The rule was stated by Henry Bent as follows: "Atomic s character concentrates in orbitals directed toward electropositive substituents".
The chemical structure of a molecule is intimately related to its properties and reactivity.
Valence bond theory proposes that molecular structures are due to covalent bonds between the atoms and that each bond consists of two overlapping and typically hybridised atomic orbitals.
Traditionally, p-block elements in molecules are assumed to hybridise strictly as sp"n", where "n" is either 1, 2, or 3.
In addition, the hybrid orbitals are all assumed to be equivalent (i.e.
the sp"n" orbitals have the same p character).
Results from this approach are usually good, but they can be improved upon by allowing hybridised orbitals with noninteger and unequal p character.
Bent's rule provides a qualitative estimate as to how these hybridised orbitals should be constructed.
Bent's rule is that in a molecule, a central atom bonded to multiple groups will hybridise so that orbitals with more s character are directed towards electropositive groups, while orbitals with more p character will be directed towards groups that are more electronegative.
By removing the assumption that all hybrid orbitals are equivalent sp"n" orbitals, better predictions and explanations of properties such as molecular geometry and bond strength can be obtained.
Bent's rule can be generalized to d-block elements as well.
The hybridisation of a metal center is arranged so that orbitals with more s character are directed towards ligands that form bonds with more covalent character.
Equivalently, orbitals with more d character are directed towards groups that form bonds of greater ionic character. 
In chemistry, pi bonds (π bonds) are covalent chemical bonds where two lobes of one involved atomic orbital overlap two lobes of the other involved atomic orbital.
Each of these atomic orbitals is zero at a shared nodal plane, passing through the two bonded nuclei.
The same plane is also a nodal plane for the molecular orbital of the pi bond.
The Greek letter π in their name refers to p orbitals, since the orbital symmetry of the pi bond is the same as that of the p orbital when seen down the bond axis.
P orbitals often engage in this sort of bonding.
D orbitals also engage in pi bonding, and form part of the basis for metal-metal multiple bonding.
Pi bonds are usually weaker than sigma bonds; the C-C double bond has a bond energy less than twice the C-C single bond bond energy; which leads to the conclusion that the p orbital overlap to form molecular orbitals is a weaker bond than when s orbitals overlap to form molecular orbitals.
From the perspective of quantum mechanics, this bond's weakness is explained by significantly less overlap between the component p-orbitals due to their parallel orientation.
This is contrasted by sigma bonds which form bonding orbitals directly between the nucleus of the bonding atoms, resulting in greater overlap and a strong sigma bond.
Pi bonds result from overlap of atomic orbitals that are in contact through two areas of overlap.
Pi-bonds are more diffuse bonds than the sigma bonds.
Electrons in pi bonds are sometimes referred to as pi electrons.
Molecular fragments joined by a pi bond cannot rotate about that bond without breaking the pi bond, because rotation involves destroying the parallel orientation of the constituent p orbitals.
For homonuclear diatomic molecules, bonding π molecular orbitals have only the one nodal plane passing through the bonded atoms, and no nodal planes between the bonded atoms.
The corresponding "anti"bonding, or π* ("pi-star") molecular orbital, is defined by the presence of an additional nodal plane between these two bonded atoms. 
In chemistry, molecular orbital (MO) theory is a method for determining molecular structure in which electrons are not assigned to individual bonds between atoms, but are treated as moving under the influence of the nuclei in the whole molecule.
Because electrons are the fundamental constituents of matter involved in bonding, their involvement in bonding has been studied exhaustively by chemists.
Electrons are shared among individual atoms in a molecule to form covalent chemical bonds.
Generally, up to three bonds can form between atoms in a molecule.
Single, or sigma covalent bonds result from the interaction between the nuclei of two discrete atoms; multiple bonds then can result due to the additional formation of pi bonds between overlapping orbitals of like symmetries.
Electrons in sigma bonds are located between the nuclei, while electrons in pi bonds are delocalized in regions above and below the nuclei.
The spatial and energetic properties of electrons within atoms are fixed by quantum mechanics to form orbitals that contain these electrons.
While atomic orbitals contain electrons ascribed to a single atom, molecular orbitals, which surround a number of atoms in a molecule, contain valence electrons between atoms.
Molecular orbital theory, which was proposed in the early twentieth century, revolutionized the study of bonding by approximating the positions of bonded electrons—the molecular orbitals—as Linear Combinations of Atomic Orbitals (LCAO).
These approximations are made by applying the Density Functional Theory (DFT) and Hartree–Fock (HF) models to the Schrödinger equation.MO theory applies the wavelike behavior of electrons as predicted by quantum mechanics, in that electrons no longer deterministically are given defined coordinates, but rather are given probable locations according to the mathematical wavefunctions defining all the possible positions of the electrons.
These wavefunctions, or electron eigenstates, quantitatively describe the atomic orbital basis in which an electron temporarily can reside.
Molecular orbitals result from the mixing of these atomic orbitals.
In this theory, each molecule has a set of molecular orbitals, in which it is assumed that the molecular orbital wave function "ψj" can be written as a simple weighted sum of the n constituent atomic orbitals "χi", according to the following equation: formula_1  
Chemical bonds between atoms in a molecule form because they make the situation more stable for the involved atoms, which generally means the sum energy level for the involved atoms in the molecule is lower than if the atoms were not so bonded.
As separate atoms approach each other to covalently bond, their orbitals affect each other's energy levels to form bonding and antibonding molecular orbitals.
The energy level of the bonding orbitals is lower, and the energy level of the antibonding orbitals is higher.
For the bond in the molecule to be stable, the covalent bonding electrons occupy the lower energy bonding orbital, which may be signified by such symbols as σ or π depending on the situation.
Corresponding anti-bonding orbitals can be signified by adding an asterisk to get σ* or π* orbitals.
A non-bonding orbital in a molecule is an orbital with electrons in outer shells which do not participate in bonding and its energy level is the same as that of the constituent atom.
Such orbitals can be designated as n orbitals.
The electrons in an n orbital are typically lone pairs.
In polyatomic molecules, different vibrational and rotational energy levels are also involved.
Roughly speaking, a molecular energy state, i.e.
an eigenstate of the molecular Hamiltonian, is the sum of the electronic, vibrational, rotational, nuclear, and translational components, such that: formula_9 where is an eigenvalue of the electronic molecular Hamiltonian (the value of the potential energy surface) at the equilibrium geometry of the molecule.
The molecular energy levels are labelled by the molecular term symbols.
The specific energies of these components vary with the specific energy state and the substance.
In molecular physics and quantum chemistry, an energy level is a quantized energy of a bound quantum mechanical state.