8.2 Hybrid Atomic Orbitals

One way to explain how chemical bonds form in diatomic molecules is by thinking in terms of overlap of atomic orbitals.

To understand how stable bonds are formed, we need a more detailed model.
The water molecule has one oxygen atom bonding to two hydrogen atoms.
The bond angle is not 90deg as shown by the evidence.
The real-world observations of a water molecule do not match the predictions of the valence bond theory model.

This isn't consistent with the evidence.

Waves combine to create new mathematical descriptions that have different shapes when atoms are bound together.

There are four equivalent hybrid orbitals that point toward the corners of a tetrahedron in the oxygen atom.

The bond angle should be a result of the overlap of the O and H orbitals.
Valence bond theory must include a hybridization component to give accurate predictions because of the observed angle of 104.5deg.

In order to make the geometry easier to see, some orbitals may be drawn in a balloon shape instead of a lump.

The experimental structure is consistent with this description.

There are no hybrid orbitals in isolated atoms.

They are formed from bonds of covalently bonding atoms.

The shapes and orientations of hybrid orbitals are very different from the atomic ones.

A set of hybrid orbitals is created.

The number of hybrid orbitals in a set is the same as the number of atomic orbitals in the set.

A set of hybrid orbitals are equivalent in shape and energy.

The type of hybrid orbitals depends on the electron-pair geometry of the atom.

S bonds are formed when hybrid orbitals overlap.

P bonds are formed by un hybridized orbitals.

We will discuss the common types of hybrid orbitals in the following sections.

A central atom with no lone pairs of electrons in a linear arrangement of three atoms is an example.

There are two regions of electron density in the molecule.
Two of the Be atom's four orbitals will mix to accommodate the two electron domains.
The number of hybrid orbitals is always equal to the number of atomic orbitals.
The half-filled hybrid orbitals will overlap with the chlorine atoms to form two identical s bonds.

The hybrid orbital is oriented in one direction.

We show the electronic differences in an isolated Be atom and in a bonded Be atom in a diagram.

The diagram has energy at the top.
One upward arrow is used to indicate one electron in an orbital and two upward arrows are used to indicate two electrons of opposite spin.

The newly created orbitals are occupied by the valence electrons.

The unpaired electron on a chlorine atom pairs up with the hybrid electron on a chlorine atom when they overlap.

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Other examples include the mercury atom in the linear HgCl2 molecule, the zinc atom in Zn(CH3)2, which has a linear C-Zn-C arrangement, and the carbon atoms in HCCH and CO2.

The University of Wisconsin-Oshkosh has orbitals in three dimensions.

There are 2 hybridized orbitals with respect to each other.

We will use these representations when the view is too crowded to see.

2 hybrid orbitals are sometimes used in crowded figures.

The molecule is trigonal and has three bonds to hydrogen atoms.

BH3 has a trigonal structure.

Three s bonds are formed in BH3 by the overlap of the three half-filled hybrid orbitals.

This includes the molecule with a lone pair on the central atom, as well as the molecule with two single bonds and a double bond connected to the central atom, such as the molecule with two single bonds and a double bond connected to the central atom.

The electrons that point in one direction are contained in a region of electron density.

There are 3 hybridized orbitals with respect to each other.

A molecule of methane, CH4, consists of a carbon atom surrounded by four hydrogen atoms.

Each carbon electron pairs with a hydrogen electron when the C-H bonds form, and the four valence electrons of the carbon atom are distributed equally in the hybrid orbitals.

A C-H s bond is created by overlap of each of the hybrid orbitals.

A sigma bond is formed by 3 orbitals of the carbon atom.

The formation of four strong, equivalent covalent bonds between the carbon atom and each of the hydrogen atoms resulted in the creation of the methane molecule, CH4.

A s bond is formed between the two carbon atoms.

The orientation of the two CH3 groups is not fixed relative to each other.
There is evidence that shows rotation around bonds.

A single pair of electrons can be held by 3 hybrid orbitals.

The nitrogen atom in ammonia is surrounded by three bonding pairs and a single pair of electrons.
3 hybridized with one hybrid.

There are two lone pairs and two bonding pairs of electrons in the water.

Two of the hybrid orbitals were occupied by lone pairs and two by bonding pairs.
Since lone pairs occupy more space than bonding pairs, structures that contain lone pairs have slightly distorted bond angles.
The observed angles in ammonia and water are slightly different.

There are five P-Cl bonds in a molecule of PCl5 and they are directed toward the corners of a trigonal bipyramid.

There are two lone pairs of electrons on the central atom, and two lone pairs on the T-shape shown.

A molecule of sulfur has six bonding pairs of electrons and a single sulfur atom.

The central atom has no lone pairs of electrons.
Two hybrid orbitals directed towards different corners of an Octahedron.

The structure around sulfur is made up of 2 orbitals.

The minor part of each orbital is not shown for clarity.

The number of regions of electron density surrounding an atom is used to determine the atom's hybridization.

The arrangements are the same as those predicted by the theory.

There are two theories that give an explanation for how the shapes of molecule are formed.

Determine the structure of the molecule.

The number of regions of electron density around an atom is determined by the number of bonds, radicals, and lone pairs that count as one region.

The set of hybridized orbitals should be assigned to this geometry.

The shapes of hybridized orbital sets are the same as the electron-pair geometries.

There are 2 orbitals arranged in a trigonal fashion.

It is important to remember that hybridization was created to rationalize observed geometries.

The model works well for molecules with small central atoms, in which the valence electron pairs are close together.
There are fewer repulsions for larger central atoms.
They do not need hybridized orbitals to explain the observed data because their compounds exhibit structures that are not consistent with the theory.
Sulfur and H2S have the same Lewis structure.
It has a smaller bond angle, which shows less hybridization on sulfur than oxygen.
There is a need to explain the structures.

Ammonium sulfate is important.

There are four regions of electron density in sulfate.

Nitrogen is sometimes used as a source of urea.