A calculation of H2S using a bond theory yields bond energies, bond lengths, and bond angles.
In our qualitative treatment, we show how overlap leads to bonding and make a rough sketch of the molecule based on the overlap.
The bond angle in H2S is 92 degrees.
In the case of H2S, a simple bond treatment matches well with the experimentally measured bond angle, which is less than 109.5deg.
We could try to explain the bonding between hydrogen and carbon using the same approach.
Only two bonds with two hydrogen atoms should be formed by carbon.
From experiments, we know that the stable compound formed from car H Bon and hydrogen is CH4 (methane), which has bond angles of 109.5deg.
A chemical bond is the overlap of two orbitals that are a product of two electrons.
The higher the overlap, the stronger the bond and the lower the mixing of two or more standard atomic orbitals.
The concentration of the electron probability density in a single direction allows for more overlap with the orbitals of other atoms.
In most cases, hybridization costs some energy.
The energy payback through bond formation is large is what causes hybridization.
The more bonds that an atom forms, the more it tends to hybridize.
The central or interior atoms form the most bonds.
The most bonds have the least tendency to hybridize.
It is important that carbon is hybridized because it tends to form four bonds in its compounds in different degrees.
There is a total number of orbitals.
Since actual energy calculations are beyond the scope of this book, we use electron geometries to predict the type of hybridization.
The hybrid orbitals all have the same energy.
The four hybrid orbitals are arranged in a way that has 109.5 degree angles between them.
Carbon's four electrons occupy the orbitals with parallel spins.
This agrees with the measured geometry of CH4.
Chemical bonds are formed because hybridized orbitals maximize overlap.
If the central atom of a molecule has lone pairs, hybrid orbitals can accommodate them.
Three of the hybrid are involved in bonding with three hydrogen atoms.
The lone pair lowers the tendency of nitrogen's orbitals to hybridize.
For a given molecule, it is beyond the scope of this text.
There are three hybrid orbitals with a trigonal geometry.
Consider H2CO if you have 2 hybrid orbitals.
Half of the orbitals are filled.
The carbon atom has four half-filled orbitals and can form four bonds: two with two hydrogen atoms and two with the oxygen atom.
The result is a pi bond when the orbitals overlap.
Only one sigma bond can be formed by two atoms.
Even though we represent the two electrons in a p bond as two half-arrows in the upper lobe, they are spread out over both the upper and lower lobes, which is one of the limitations we encounter when we try to represent electrons with arrows.
We can label the bonds in the molecule using a method that tells us the type of bond as well as the type of overlap.
Two single bonds and one double bond are formed by the central carbon atom.
The theory of the bonds gives us more insight.
The double bond between carbon and oxygen consists of two different kinds of bonds--one s and one p--whereas in the Lewis model the two bonds within the double bond appear identical.
P bonds are weaker than s bonds because side-to-side overlap is less efficient than end-to-end One.
The additional bonds must be the same as the s bonds.
Valence bond theory gives us more insight into the bonds.
The types of orbitals involved in bonding are shown in the valence bond theory.
The resulting geometry of the molecule is predicted by the 2 hybrid orbitals on the central atom.
The HCO bond angle is close to the predicted values and the HCH bond angle is close to the measured angle.
The rotation of a double bond is severely restricted by virtue bond theory.
rotation about a single bond is relatively unrestricted, even though it is highly restricted.
The rotation of the sigma single bond was relatively free.
Light is detected by a chemical switch in the human eye.
The sensitive cells called rods and cones are the s bond.
retinal has about that bond.
retinal causes changes in theprotein to which it is bound.
There are two different forms of 1,2-dichloroethene.
Nature can make different compounds out of the same atoms by arranging them in different ways.
In Chapter 22 we will discussomerism and how important it is in organic chemistry.
Double bonds are stronger and shorter than single bonds.
A double bond is composed of two different bonds, one s and one p.
A double bond is the sharing of two electron pairs.
The C " C double bond is not an exception.
A double bond is more strong than a single bond.
Each carbon atom has four half-filled orbitals and can form four bonds: one with a hydrogen atom and three with the other carbon atom.
According to the Lewis model, elements occurring in the third row of the periodic table can exhibit expanded octets.
There are 2 hybrid orbitals.
Consider sulfur hexafluoride, SF6.
The geometry of the sulfur atom is in agreement with the theory and observed geometry.
Examples of the five main types of atomic orbital hybridization have been studied.
Computational valence bond theory uses a computer to calculate the energy of the molecule, with the degree of hybridization as well as the type of hybridization being varied to find the combination that gives the molecule the lowest overall energy.
The table shows the five VSEPR electron geometries and the corresponding hybridization schemes.
When H2S is largely un hybridized, it is the best we can do.
The Lewis model and valence bond theory are going to be put together to describe bonding.
This involves drawing the Lewis structure for the molecule, determining its geometry using VSEPR theory, determining the hybridization of the interior atoms, and labeling each bond with the s and p notation.
This procedure involves everything you have learned about bonding in this chapter and chapter 10.
Two examples of how to apply the procedure in the columns to the right are shown in the left column.