11 -- Part 2: Chemical Bonding II: Valence Bond and Molecular Orbital Theories
It's not appropriate to describe bond formation in terms of hybrid orbitals.
In organic chemistry, the concept of hybridization is used a lot.
The central atoms in H2O and NH3 need an sp3 hybridization scheme.
The bond angles in water and NH3 are similar to these angles.
Bonding in NH3 can be described in terms of the following diagram.
The three half-filled sp3 orbitals are involved in bond formation.
The observed bond angles do not conform to the expected H bond angle.
H bonds to open up to hybrid orbitals.
There is a single theory that is consistent with all the available evidence.
The sp2 hybridization scheme is related to the trigonal-planar electron group geometry.
Bonding in a linear arrangement of atoms is described by the sp hybrid orbitals.
When this scheme is used to describe bonds formed by atoms other than beryllium, there are two un hybridized p orbitals that are oriented along the y and Z axis.
Themolecular are left un hybridized.
The scheme is further outlined in theory, and the diagrams of the valence-shell are shown here.
CONCEPT ASSESSMENT can be used to form bonds.
The text states that orbital mixing is a mathematical process of changing pure atomic orbitals for isolated atoms into new atomic orbitals for bonding atoms.
A hybrid atomic orbital is a result of a mathematical combination of wave functions describing two or more atomic orbitals.
The linear combinations of the s and p are depicted.
The three-dimensional forms of the sp hybrid are shown above the maps.
We have previously said that hybridization is not a real phenomenon, but an after-the-fact rationalization of an experiment.
There is no better example of this point than the issue of the sp3d and sp3d2 hybrid orbitals.
In the previous section, we used either the experimental geometry or the geometry predicted by the VSEPR theory to help us decide on the appropriate hybridization scheme for the central atom.
We used the concept of hybridization in Chapter 10.
The known geometries of CH4, H2O, and NH3 are accounted for by the sp3d and sp3d2 hybrid orbitals.
The same number of hybrid orbitals should be produced by the same scheme as there is by the central atom.
An sp3 hybridization scheme for the central atom predicts that four hybrid orbitals are distributed in a tetrahedral fashion.
Depending on how many hybrid orbitals are involved in orbital overlap and how many contain lone-pair electrons, the structures are either trigonal-pyramidal or angular.
Some important applications to organic chemistry will be considered in the next section.
Predict the shape of the XeF4 molecule with a hybridization scheme consistent with this prediction.
A plausible Lewis structure is what you should write.
The Xe atom's valence shell must be expanded to accommodate 12 electrons in Chemical Bonding II: Valence Bond and Orbital Lewis structure.
The central atom has an electron-group geometry.
Bond pairs and lone pairs are the electron groups.
The four pairs of bond electrons are directed to the corners of a square, and the lone pairs of electrons are found above and below the plane of the Xe and F atoms.
Pick a scheme that matches the prediction.
There are only a few pairs of electrons above and below the plane of the Xe and F atoms.
The number of electron pairs in the geometry of the electron-group determines how many orbitals are used in the scheme.
The shape of and bonding in a molecule can be described with a combination of VSEPR and hybridization theory.
The sp3d2 hybridization scheme for xenon is not the best way to describe the bonding in XeF.
There is a central atom in the ion Cl2F+.
The central atom in the ion BrF + 4 is the subject of a possible hybridization scheme.
There isn't a correct method for describing structures.
The experimental evidence from which the structure is established is the only correct information.
You may be able to rationalize the evidence by using one method or another once the evidence is in hand.
The valence bond method seems to do a better job of explaining the observed 92deg bond angle than the VSEPR theory.
The method used by the VSEPR theory suggests a 109.5 degree angle to the electron-group geometry.
The predicted bond angle is less than geometry because of the data that was changed to accommodate lone-pair-lone-pair and lone approximate molecular pair-bond-pair repulsions.
Unless you have specific information to suggest otherwise, describing a molecule in bonding based on shape is a good bet.
It's important to remember that both geometry.
There are two different types of overlap when multiple bonds are described.
Specific examples of the carbon-to-carbon double bond in ethylene, C2H4 and the carbon-to-carbon triple bond in acetylene will be used in our discussion.
There is a carbon-to-carbon double bond in the Lewis structure.
There is a molecule called ethylene.
The bond angles are close.
The CH2 groups are coplanar.
sp2 is the scheme that produces a set of hybrid orbitals.
The line joining the two atoms overlaps.
Above and below the plane of the carbon and hydrogen atoms, there is a region of high electron density.
A p bond is formed by side-to-side fashion.
The phase of the p orbitals is retained.
An alternative definition for s and p bonds is based on the number of nodal planes that are parallel to the bond axis.
CH2 p bond in a double bond and group out of the plane of the other would reduce overlap of two in a triple bond.
The C2H4 molecule is planar and the double bond is rigid.
The s bond involves more overlap than the p bond.
The bond angles are 120 degrees.
One of the carbon-to-carbon bonds is a s bond.
One p bond is formed by a pair of p orbitals, which are oriented along an axis in the plane of the page.
The lower part of the diagram shows the formation of the other p bond.
One of the p orbitals has a transparent surface to distinguish it from the other p bond.
The number of electron groups around a central atom affects the number of atomic orbitals.
The carbon atom in formaldehyde contains three electron groups, which means that three atomic orbitals are hybridized to form three sp2 orbitals.
The central atom is C, and the terminal atoms are H and O.
There are 12 valence electrons.
This structure requires a carbon-to-oxygen double bond.
The central C atom has an electron-group geometry.
The central C atom is the basis of the s-bond framework.
The distribution of three electron groups suggests a trigonal-planar molecule.
The scheme that conforms to the electron-group geometry is identified.
sp2 hybrid orbitals are associated with a trigonal-planar orientation.
The orbital set sp2 + p is produced by the hybridized C atom.
Two sp2 hybrid orbitals are used to form bonds with H atoms.
A s bond with oxygen is formed by the remaining sp2 hybrid orbital.
The central and terminal atoms have bonding orbitals.
The bonds are used to bond two H atoms and one O atom.
We look at the number of bonds and lone pairs around the atom to determine the electron-group geometry and bond angles.
There is a plausible bonding scheme for dimethyl ether.
It is difficult to draw three-dimensional sketches to show overlaps.
Straight lines are used to draw bonds between atoms.
They are labeled s or p, and the orbitals that overlap are indicated.
A Lewis structure is the first step in describing a bonding scheme.
A description of the species obtained by experiment is sometimes the starting point.
A formula with bond angles is given.
The scheme should be consistent with this structure.
The angle of the figure is close to the angle of the book.
An sp3 hybridization scheme is employed by the O atom on the Bonding and structure of right.
The lone pair electrons are not shown in the original structure.
The observed bond angles can be used to infer the presence of lone pairs.
The sp3 hybridization at the oxygen atom would have been deduced using a Lewis structure and VSEPR theory.
An industrial solvent is acetonitrile.
The scheme should be consistent with its structure.
The molecule diazine has a formula.
The combination of Lewis structures, VSEPR theory, and the valence bond method makes for a potent description of bonds.
Most of the time, they are satisfactory.
Sometimes a greater understanding of structures and properties is needed by chemists.
Oxygen is paramagnetic, H + 2 is a stable species, and the electronic spectrum of molecule is not explained by any of these methods.
A different method of describing chemical bonding is needed to address these questions.
We will only give an overview and focus on the application of the theory to the diatomic molecule.
We will start our overview by comparing the conceptual model we introduced in Chapter 8 for multielectron atoms with the molecular orbital theory.
In a similar way, we imagine that each molecule has a set of orbitals and can be built up by placing electrons into them.
Unlike atomic orbitals, which are centered on a single nucleus,molecular orbitals are defined by all the nuclei.
Let's look at the results of a calculation for F in more detail.
The equation for the F molecule is used to solve the total wave function.
The idea is that because a molecule is composed of atoms, the AOs of those atoms can be used as the basis of a method for describing how each electron in a molecule interacts with all nuclei simultaneously.
The atomic orbitals can be used as a basis for describing a molecule.
The premise of the LCAO method is that each MO can be represented as a linear combination of all AOs.
The following ideas are incorporated into the LCAO method.
When the atoms combine to form a molecule, the AOs are replaced by a set of molecular orbitals.
This conceptualization applies what we know about electrons in atoms and extends that description to Molecules.
Deducing the appropriate combinations of AOs is a difficult aspect of applying the approach.
The appropriate combinations are simple for diatomic molecules.
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