Molecular Orbital Theory The goal of molecular orbital theory is to describe molecules in a similar way to how we describe atoms, that is, in terms of orbitals, orbital diagrams, and electron configurations. For example, to give you a glimpse at where we are headed, the following are orbital diagrams for O2 and O.

O2 O
Each line in the molecular orbital diagram represents a molecular orbital, which is the volume within which a high percentage of the negative charge generated by the electron is found. The molecular orbital volume encompasses the whole molecule. We assume that the electrons would fill the molecular orbitals of molecules like electrons fill atomic orbitals in atoms.
The molecular orbitals are filled in a way that yields the lowest potential energy for the molecule.
The maximum number of electrons in each molecular orbital is two. (We follow the Pauli exclusion principle.)
Orbitals of equal energy are half filled with parallel spin before they begin to pair up. (We follow Hund's Rule.)
Before we continue with a description of a model used to generate molecular orbital diagrams, lets get a review of light and electron waves and how two waves can interact. The wave description of light describes the effect that the light has on the space around it. This effect is to generate an oscillating electric and magnetic fields. These fields can vary in intensity, which is reflected in varying brightness of light.

The wave description of the electron describes the variation in the intensity of negative charge generated by the electron.

Light waves can interact in-phase, which leads to an increase in the intensity of the light (brighter) and out-of-phase, which leads to a decrease in the intensity of the light (less bright). Electron waves can also interact in-phase and out-of-phase. In-phase interaction leads to an increase in the intensity of the negative charge. Out-of-phase interaction leads to a decrease in the intensity of the negative charge.
One common approximation that allows us to generate molecular orbital diagrams for some small diatomic molecules is called the Linear Combination of Atomic Orbitals (LCAO) approach. The following assumptions lie at the core of this model.
Molecular orbitals are formed from the overlap of atomic orbitals.
Only atomic orbitals of about the same energy interact to a significant degree.
When two atomic orbitals overlap, they interact in two extreme ways to form two molecular orbitals, a bonding molecular orbital and an antibonding molecular orbital.
For example, our model assumes that two 1s atomic orbitals can overlap in two extreme ways to form two molecular orbitals. One of the ways the atomic orbitals interact is in-phase, which leads to wave enhancement similar to the enhancement of two in-phase light waves. Where the atomic orbitals overlap, the in-phase interaction leads to an increase in the intensity of the negative charge in the region where they overlap. This creates an increase in negative charge between the nuclei and an increase in the plus-minus attraction between the electron charge and the nuclei for the atoms in the bond. The greater attraction leads to lower potential energy. Because electrons in the molecular orbital are lower potential energy than in separate atomic orbitals, energy would be required to shift the electrons back into the 1s orbitals of separate atoms. This keeps the atoms together in the molecule, so we call this orbital a bonding molecular orbital. The molecular orbital formed is symmetrical about the axis of the bond. Symmetrical molecular orbitals are called sigma, σ, molecular orbitals. The symbol σ1s is used to describe the bonding molecular orbital formed from two 1s atomic orbitals.
The second way that two atomic orbitals interact is out-of-phase. Where the atomic orbitals overlap, the out-of-phase interaction leads to a decrease in the intensity of the negative charge. This creates a decrease in negative charge between the nuclei and a decrease in the plus-minus attraction between the electron charge and the nuclei for the atoms in the bond. The lesser attraction leads to higher potential energy. The electrons are more stable in the 1s atomic orbitals of separate atoms, so electrons in this type of molecular orbital destabilize the bond between atoms. We call molecular orbitals of this type antibonding molecular orbitals. The molecular orbital formed is symmetrical about the axis of the bond, so it is a sigma molecular orbital with a symbol of σ*1s. The asterisk indicates an antibonding molecular orbital.
The following diagram shows the bonding and antibonding molecular orbitals formed from the interaction of two 1s atomic orbitals.


When two larger atoms atoms combine to form a diatomic molecule (like O2, F2, or Ne2), more atomic orbitals interact. The LCAO approximation assumes that only the atomic orbitals of about the same energy interact. For O2, F2, or Ne2, the orbital energies are different enough so only orbitals of the same energy interact to a significant degree.
Like for hydrogen, the 1s from one atom overlaps the 1s from the other atom to form a σ1s bonding molecular orbital and a σ*1s antibonding molecular orbital. The shapes would be similar to those formed from the 1s orbitals for hydrogen. The 2s atomic orbital from one atom overlaps the 2s from the other atom to form a σ2s bonding molecular orbital and a σ*2s antibonding molecular orbital. The shapes of these molecular orbitals would be similar to those for the σ1s and σ*1s molecular orbitals. Both σ2s and σ*2s molecular orbitals are higher energy and larger than the σ1s and σ*1s molecular orbitals.
The p atomic orbitals of the two atoms can interact in two different ways, parallel or end-on. The molecular orbitals are different for each type of interaction. The end-on interaction between two 2px atomic orbitals yields sigma molecular orbitals, which are symmetrical about the axis of the bond.

The two 2py atomic orbitals overlap in parallel and form two pi molecular orbitals. Pi molecular orbitals are asymmetrical about the axis of the bond.

The 2pz-2pz overlap generates another pair of π2p and π*2p molecular orbitals. The 2pz-2pz overlap is similar to the The 2py-2py overlap. To visualize this overlap, picture all of the orbitals in the image above rotated 90 degrees so the axes that run through the atomic and molecular orbitals are perpendicular to the screen (paper). The molecular orbitals formed have the same potential energies as the molecular orbitals formed from the 2py-2py overlap.
There is less overlap for the parallel atomic orbitals. When the interaction is in-phase, less overlap leads to less electron charge enhancement between the nuclei. This leads to less electron charge between the nuclei for the pi bonding molecular orbital than for the sigma bonding molecular orbital. Less electron character between the nuclei means less plus-minus attraction, less stabilization, and higher potential energy for the pi bonding molecular orbital compared to the sigma bonding molecular orbital.
When the interaction is out-of-phase, less overlap leads to less shift of electron charge from between the nuclei. This leads to more electron charge between the nuclei for the pi antibonding molecular orbital than for the sigma antibonding molecular orbital. More electron charge between the nuclei means more plus-minus attraction and lower potential energy for the pi antibonding molecular orbital compared to the sigma antibonding molecular orbital.
The expected molecular orbital diagram from the overlap of 1s, 2s and 2p atomic orbitals is as follows. We will use this diagram to describe O2, F2, Ne2, CO, and NO.

We use the following procedure when drawing molecular orbital diagrams.
Determine the number of electrons in the molecule. We get the number of electrons per atom from their atomic number on the periodic table. (Remember to determine the total number of electrons, not just the valence electrons.)
Fill the molecular orbitals from bottom to top until all the electrons are added. Describe the electrons with arrows. Put two arrows in each molecular orbital, with the first arrow pointing up and the second pointing down.
Orbitals of equal energy are half filled with parallel spin before they begin to pair up.
We describe the stability of the molecule with bond order.
bond order = 1/2 (#e- in bonding MO's - #e- in antibonding MO's)
We use bond orders to predict the stability of molecules.
If the bond order for a molecule is equal to zero, the molecule is unstable.
A bond order of greater than zero suggests a stable molecule.
The higher the bond order is, the more stable the bond.
We can use the molecular orbital diagram to predict whether the molecule is paramagnetic or diamagnetic. If all the electrons are paired, the molecule is diamagnetic. If one or more electrons are unpaired, the molecule is paramagnetic.
EXAMPLES:
1. The molecular orbital diagram for a diatomic hydrogen molecule, H2, is

The bond order is 1. Bond Order = 1/2(2 - 0) = 1
The bond order above zero suggests that H2 is stable.
Because there are no unpaired electrons, H2 is diamagnetic.
2. The molecular orbital diagram for a diatomic helium molecule, He2, shows the following.

The bond order is 0 for He2. Bond Order = 1/2(2 - 2) = 0
The zero bond order for He2 suggests that He2 is unstable.
If He2 did form, it would be diamagnetic.
3. The molecular orbital diagram for a diatomic oxygen molecule, O2, is

O2 has a bond order of 2. Bond Order = 1/2(10 - 6) = 2
The bond order of two suggests that the oxygen molecule is stable.
The two unpaired electrons show that O2 is paramagnetic.
4. The molecular orbital diagram for a diatomic fluorine molecule, F2, is

F2 has a bond order of 1. Bond Order = 1/2(10 - 8) = 1
The bond order of one suggests that the fluorine molecule is stable.
Because all of the electrons are paired, F2 is diamagnetic.
5. The molecular orbital diagram for a diatomic neon molecule, Ne2, is

Ne2 has a bond order of 0. Bond Order = 1/2(10 - 10) = 0
The zero bond order for Ne2 suggests that Ne2 is unstable.
If Ne2 did form, it would be diamagnetic.
We can describe diatomic molecules composed of atoms of different elements in a similar way. The bond between the carbon and oxygen in carbon monoxide is very strong despite what looks like a strange and perhaps unstable Lewis Structure.

The plus formal charge on the more electronegative oxygen and the minus formal charge on the less electronegative carbon would suggest instability. The molecular orbital diagram predicts CO to be very stable with a bond order of three.







We predict the nitrogen monoxide molecule to be unstable according to the Lewis approach to bonding.

The unpaired electron and the lack of an octet of electrons around nitrogen would suggest an unstable molecule. NO is actually quite stable. The molecular orbital diagram predicts this by showing the molecule to have a bond order of 2.5.

Any precipitation, including snow, that contains a heavy concentration of sulfuric and nitric acids. This form of pollution is a serious environmental problem in the large urban and industrial areas of North America, Europe, and Asia. Automobiles, certain industrial operations, and electric power plants that burn fossil fuels emit the gases sulfur dioxide and nitrogen oxide into the atmosphere, where they combine with water vapour in clouds to form sulfuric and nitric acids. The highly acidic precipitation from these clouds may contaminate lakes and streams, damaging fish and other aquatic species; damage vegetation, including agricultural crops and trees; and corrode the outsides of buildings and other structures (historic monuments are especially vulnerable). Though usually most severe around large urban and industrial areas, acid precipitation may also occur at great distances from the source of the pollutants

The presence in the atmospheric environment of natural and artificial substances that affect human health or well-being, or the well-being of any other specific organism. Pragmatically, air pollution also applies to situations where contaminants impact structures and artifacts or esthetic sensibilities (such as visibility or smell). Most artificial impurities are injected into the atmosphere at or near the Earth's surface.

SOURCES


Different types of pollution are conveniently specified in various ways: gaseous, such as carbon monoxide, or particulate, such as smoke, pesticides, and aerosol sprays; inorganic, such as hydrogen fluoride, or organic, such as mercaptans; oxidizing substances, such as ozone, or reducing substances, such as oxides of sulfur and oxide s of nitrogen; radioactive substances, such as iodine-131; inert substances, such as pollen or fly ash; or thermal pollution, such as the heat produced by nuclear power plants.

Air contaminants are produced in many ways and come from many sources; it is difficult to identify all the various producers. Also, for some pollutants such as carbon dioxide and methane, the natural emissions sometimes far exceed the anthropogenic emissions.

Both anthropogenic and natural emissions are variable from year to year, depending on fuel usage, industrial development, and climate. In some countries where pollution control regulations have been implemented, emissions have been significantly reduced. For example, in the United States sulfur dioxide emissions dropped by about 30% between 1970 and 1992, and carbon monoxide (CO) emissions were cut by over 30% in the same period. However, in some developing countries emissions continually rise as more cars are put on the road and more industrial facilities and power plants are constructed. In dry regions, natural emissions of nitrogen oxides (NOx), carbon dioxide (CO2), and hydrocarbons can be greatly increased during a season with high rainfall and above-average vegetation growth.

The anthropogenic component of most estimates of the methane budget is about two-thirds. Ruminant production and emissions from rice paddies are regarded as anthropogenic because they result from human agricultural activities. The perturbations to carbon dioxide since the industrial revolution are also principally the result of human activities. These emissions have not yet equilibrated with the rest of the carbon cycle and so have had a profound effect on atmospheric levels, even though emissions from fossil fuel combustion are dwarfed by natural emissions.

EFFECTS

The major concern with air pollution relates to its effects on humans. Since most people spend most of their time indoors, there has been increased interest in air-pollution concentrations in homes, workplaces, and shopping areas. Much of the early information on health effects came from occupational health studies completed prior to the implementation of general air-quality standards.

Air pollution principally injures the respiratory system, and health effects can be studied through three approaches, clinical, epidemiological, and toxicological. Clinical studies use human subjects in controlled laboratory conditions, epidemiological studies assess human subjects (health records) in real-world conditions, and toxicological studies are conducted on animals or simple cellular systems. Of course, epidemiological studies are the most closely related to actual conditions, but they are the most difficult to interpret because of the lack of control and the subsequent problems with statistical analysis. Another difficulty arises because of differences in response among different people. For example, elderly asthmatics are likely to be more strongly affected by sulfur dioxide than the teenage members of a hiking club. See also Epidemiology.

Damage to vegetation by air pollution is of many kinds. Sulfur dioxide may damage field crops such as alfalfa and trees such as pines, especially during the growing season (Fig. 1). Both hydrogen fluoride (HF) and nitrogen dioxide (NO2) in high concentrations have been shown to be harmful to citrus trees and ornamental plants, which are of economic, importance in central Florida. Ozone and ethylene are other contaminants that cause damage to certain kinds of vegetation.

Air pollution can affect the dynamics of the atmosphere through changes in longwave and shortwave radiation processes. Particles can absorb or reflect incoming short-wave solar radiation, keeping it from the Earth's surface during the day. Greenhouse gases can absorb long-wave radiation emitted by the Earth's surface and atmosphere.

Carbon dioxide, methane, fluorocarbons, nitrous oxides, ozone, and water vapor are important greenhouse gases. These represent a class of gases that selectively absorb long-wave radiation. This effect warms the temperature of the Earth's atmosphere and surface higher than would be found in the absence of an atmosphere (the greenhouse effect). Because the amount of greenhouse gases in the atmosphere is rising, there is a possibility that the temperature of the atmosphere will gradually rise, possibly resulting in a general warming of the global climate over a time period of several generations. See also Greenhouse effect.

Researchers are also concerned with pollution of the stratosphere (10–50 km or 6–30 mi above the Earth's surface) by aircraft and by broad surface sources. The stratosphere is important, because it contains the ozone layer, which absorbs part of the Sun's short-wave radiation and keeps it from reaching the surface. If the ozone layer is significantly depleted, an increase in skin cancer in humans is expected. Each 1% loss of ozone is estimated to increase the skin cancer rate 3–6%. See also Stratosphere.

Visibility is reduced as concentrations of aerosols or particles increase. The particles do not just affect visibility by themselves but also act as condensation nuclei for cloud or haze formation. In each of the three serious air-pollution episodes discussed above, smog (smoke and fog) were present with greatly reduced visibility.

CHEMISTRY

Air pollution can be divided into primary and secondary compounds, where primary pollutants are emitted directly from sources (for example, carbon monoxide, sulfur dioxide) and secondary pollutants are produced by chemical reactions between other pollutants and atmospheric gases and particles (for example, sulfates, ozone). Most of the chemical transformations are best described as oxidation processes. In many cases these secondary pollutants can have significant environmental effects, such as acid rain and smog.

Smog is the best-known example of secondary pollutants formed by photochemical processes, as a result of primary emissions of nitric oxide (NO) and reactive hydrocarbons from anthropogenic sources such as transportation and industry as well as natural sources. Energy from the Sun causes the formation of nitrogen dioxide, ozone (O3), and peroxyacetalnitrate, which cause eye irritation and plant damage.

It has been shown that when emissions of sulfur dioxide and nitrogen oxide from tall power plant and other industrial stacks are carried over great distances and combined with emissions from other areas, acidic compounds can be formed by complex chemical reactions. In the absence of anthropogenic pollution sources, the average pH of rain is around 5.6 (slightly acidic). In the eastern United States, acid rain with a pH less than 5.0 has been measured and consists of about 65% dilute sulfuric acid, 30% dilute nitric acid, and 5% other acids.

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