However, their direct emissions are typically lower than those of comparable conventional vehicles. Well-to-wheel emissions include all emissions related to fuel production, processing, distribution, and use. In the case of gasoline, emissions are produced while extracting petroleum from the earth, refining it, distributing the fuel to stations, and burning it in vehicles. In the case of electricity, most electric power plants produce emissions, and there are additional emissions associated with the extraction, processing, and distribution of the primary energy sources they use for electricity production.
More in this section Therefore, the methane molecule cannot be adequately represented by simple overlap of the 2s and 2p orbitals of carbon with the 1s orbitals of each hydrogen atom. To explain the bonding in methane, it is necessary to introduce the concept of hybridization and hybrid atomic orbitals. Hybridization is the mixing of the atomic orbitals in an atom to produce a set of hybrid orbitals. When hybridization occurs, it must do so as a result of the mixing of nonequivalent orbitals.
In other words, s and p orbitals can hybridize but p orbitals cannot hybridize with other p orbitals. Hybrid orbitals are the atomic orbitals obtained when two or more nonequivalent orbitals form the same atom combine in preparation for bond formation.
In the current case of carbon, the single 2s orbital hybridizes with the three 2p orbitals to form a set of four hybrid orbitals, called sp 3 hybrids see Figure 3 below. The sp 3 hybrids are all equivalent to one another. Spatially, the hybrid orbitals point towards the four corners of a tetrahedron. The process of sp 3 hybridization is the mixing of an s orbital with a set of three p orbitals to form a set of four sp 3 hybrid orbitals.
Each large lobe of the hybrid orbitals points to one corner of a tetrahedron. The four lobes of each of the sp 3 hybrid orbitals then overlap with the normal unhybridized 1s orbitals of each hydrogen atoms to form the tetrahedral methane molecule. Use the link below to answer the following questions.
Read only the sections on ammonia and water hybridization. Romeo and Juliet were two of the great lovers of all time. Their embrace allowed no other person to be a part of it — they only wanted to be with each other. It took outside intervention parents are like that! Paired electrons are similar to the lovers. They do not bond covalently until they are unpaired. Then they can become a part of a larger chemical structure.
The beryllium atom contains all paired electrons and so must also undergo hybridization. One of the 2s electrons is first promoted to the empty 2p x orbital see figure below. Now the hybridization takes place only with the occupied orbitals and the result is a pair of sp hybrid orbitals. The two remaining p orbitals p y and p z do not hybridize and remain unoccupied see Figure 6 below. The geometry of the sp hybrid orbitals is linear, with the lobes of the orbitals pointing in opposite directions along one axis, arbitrarily defined as the x-axis see Figure 7.
Each can bond with a 1s orbital from a hydrogen atom to form the linear BeH 2 molecule. Figure 7. The process of sp hybridization is the mixing of an s orbital with a single p orbital the pxorbital by convention , to form a set of two sp hybrids.
The two lobes of the sp hybrids point opposite one another to produce a linear molecule. Other molecules whose electron domain geometry is linear and for whom hybridization is necessary also form sp hybrid orbitals.
Examples include CO 2 and C 2 H 2 , which will be discussed in further detail later. First a paired 2s electron is promoted to the empty 2p y orbital see Figure 8. This is followed by hybridization of the three occupied orbitals to form a set of three sp 2 hybrids, leaving the 2p z orbital unhybridized see Figure 9.
Quantum-mechanical calculations suggest why the observed bond angles in H 2 O differ from those predicted by the overlap of the 1 s orbital of the hydrogen atoms with the 2 p orbitals of the oxygen atom. When atoms are bound together in a molecule, the wave functions combine to produce new mathematical descriptions that have different shapes. This process of combining the wave functions for atomic orbitals is called hybridization and is mathematically accomplished by the linear combination of atomic orbitals , LCAO, a technique that we will encounter again later.
The new orbitals that result are called hybrid orbitals. The valence orbitals in an isolated oxygen atom are a 2 s orbital and three 2 p orbitals. The valence orbitals in an oxygen atom in a water molecule differ; they consist of four equivalent hybrid orbitals that point approximately toward the corners of a tetrahedron Figure 2.
Consequently, the overlap of the O and H orbitals should result in a tetrahedral bond angle The observed angle of In the following sections, we shall discuss the common types of hybrid orbitals. The beryllium atom in a gaseous BeCl 2 molecule is an example of a central atom with no lone pairs of electrons in a linear arrangement of three atoms.
There are two regions of valence electron density in the BeCl 2 molecule that correspond to the two covalent Be—Cl bonds. This hybridization process involves mixing of the valence s orbital with one of the valence p orbitals to yield two equivalent sp hybrid orbitals that are oriented in a linear geometry Figure 3. In this figure, the set of sp orbitals appears similar in shape to the original p orbital, but there is an important difference.
The number of atomic orbitals combined always equals the number of hybrid orbitals formed. The p orbital is one orbital that can hold up to two electrons. The two electrons that were originally in the s orbital are now distributed to the two sp orbitals, which are half filled.
We illustrate the electronic differences in an isolated Be atom and in the bonded Be atom in the orbital energy-level diagram in Figure 4. These diagrams represent each orbital by a horizontal line indicating its energy and each electron by an arrow. Energy increases toward the top of the diagram. We use one upward arrow to indicate one electron in an orbital and two arrows up and down to indicate two electrons of opposite spin.
When atomic orbitals hybridize, the valence electrons occupy the newly created orbitals. The Be atom had two valence electrons, so each of the sp orbitals gets one of these electrons. Each of these electrons pairs up with the unpaired electron on a chlorine atom when a hybrid orbital and a chlorine orbital overlap during the formation of the Be—Cl bonds.
Any central atom surrounded by just two regions of valence electron density in a molecule will exhibit sp hybridization. Check out the University of Wisconsin-Oshkosh website to learn about visualizing hybrid orbitals in three dimensions.
The valence orbitals of a central atom surrounded by three regions of electron density consist of a set of three sp 2 hybrid orbitals and one unhybridized p orbital. This arrangement results from sp 2 hybridization, the mixing of one s orbital and two p orbitals to produce three identical hybrid orbitals oriented in a trigonal planar geometry Figure 5.
The observed structure of the borane molecule, BH 3, suggests sp 2 hybridization for boron in this compound. The molecule is trigonal planar, and the boron atom is involved in three bonds to hydrogen atoms Figure 7.
We can illustrate the comparison of orbitals and electron distribution in an isolated boron atom and in the bonded atom in BH 3 as shown in the orbital energy level diagram in Figure 8. We redistribute the three valence electrons of the boron atom in the three sp 2 hybrid orbitals, and each boron electron pairs with a hydrogen electron when B—H bonds form.
Any central atom surrounded by three regions of electron density will exhibit sp 2 hybridization. This includes molecules with a lone pair on the central atom, such as ClNO Figure 9 , or molecules with two single bonds and a double bond connected to the central atom, as in formaldehyde, CH 2 O, and ethene, H 2 CCH 2. The valence orbitals of an atom surrounded by a tetrahedral arrangement of bonding pairs and lone pairs consist of a set of four sp 3 hybrid orbitals. The hybrids result from the mixing of one s orbital and all three p orbitals that produces four identical sp 3 hybrid orbitals Figure Each of these hybrid orbitals points toward a different corner of a tetrahedron.
A molecule of methane, CH 4 , consists of a carbon atom surrounded by four hydrogen atoms at the corners of a tetrahedron. The carbon atom in methane exhibits sp 3 hybridization.
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