Carbon dioxide hybrid orbitals

Carbon dioxide, (a) Overlap of carbon sp hybrid orbitals with oxygen 2px orbitals to give a bonds O—C—O. (b) The two tt molecular orbitals of the molecule formed from unused 2py and 2pz orbitals. 

Self-Test 3.7B Suggest a structure in terms of hybrid orbitals for carbon dioxide, C02. 

In terms of this model, two effective pairs around an atom will always require sp hybridization of that atom. The sp orbitals of carbon in carbon dioxide are shown in Fig. 14.15, and the corresponding orbital energy-level diagram for their formation is given in Fig. 14.16. The sp hybrid orbitals are used to form the a bonds between carbon and the oxygen atoms. Note that two 2p orbitals remain unchanged on the sp hybridized carbon. They are used to form the it bonds to the oxygen atoms. 

Now we are ready to account for the Lewis structure of carbon dioxide. The sp orbitals on carbon form cr bonds with the sp2 orbitals on the two oxygen atoms (Fig. 14.15). The remaining sp2 orbitals on the oxygen atoms hold lone pairs. The v bonds between the carbon atom and each oxygen atom are formed by the overlap of parallel 2p orbitals. The sp hybridized carbon atom... 

Carbon dioxide is linear as it has only two C-0 a bonds and no lone pairs on C. The C atom must be sp hybridized and the only trick is to get the two n bonds orthogonal to each other. They must be like that because the p orbitals on C involved in the two n bonds are themselves orthogonal (2pj, and 2pz). Most people would draw the O atoms as sp hybridized, rather than sp or even unhybridized, but this is unimportant as you can t really tell. 

Allene (see Problem 1.41) is related structurally to carbon dioxide, CO2. Draw a picture showing the orbitals involved in the a and it bonds of CO2, and identify the hybridization of carbon. 

From the Lewis structure of the carbon dioxide molecule (Figure 2- 18a) it is seen that the carbon atom is surrounded by two electron groups (two double bonds). Two electron groups mean that there is a need for two identical orbitals 180° apart according to the VSEPR theory and Table 2- 1 on page 70. The carbon atom solves this problem by forming two identical so-called sp-hybrid orbital. As the name sp indicated these orbitals are made from one s-orbital and one p-orbital. 

Again we see that in double bonds we have o-bonds in the overlap between hybrid orbitals and 71-bonds in the space between atomic p-orbitals. Thus in the case of carbon dioxide the two 71-bonds are rotated 90° relative to each other. 

In Section 1.3, we saw that for molecules with one covalent bond, the dipole moment of the bond is identical to the dipole moment of the molecule. For molecules that have more than one covalent bond, the geometry of the molecule must be taken into account because both the magnitude and the direction of the individual bond dipole moments (the vector sum) determine the overall dipole moment of the molecule. Symmetrical molecules, therefore, have no dipole moment. For example, let s look at the dipole moment of carbon dioxide (CO2). Because the carbon atom is bonded to two atoms, it uses sp orbitals to form the C—O a bonds. The remaining two p orbitals on carbon form the two C—O tt bonds. The individual carbon-oxygen bond dipole moments cancel each other— because sp orbitals form a bond angle of 180°—giving carbon dioxide a dipole moment of zero D. Another symmetrical molecule is carbon tetrachloride (CCI4). The four atoms bonded to the sp hybridized carbon atom are identical and project symmetrically out from the carbon atom. Thus, as with CO2, the symmetry of the molecule causes the bond dipole moments to cancel. Methane also has no dipole moment. 

When a carbon atom is attached to two other atoms, as in acetylene or carbon dioxide, there is sp hybridization and the bonds lie in a straight line. The electronic structure of carbon dioxide is shown in Figure 1.19. Two % bonds are now formed in addition to the a bonds. In the figure, the unhybridized p orbitals above and below the plane overlap with the p orbitals of the right-hand oxygen atom, and the unhybridized p orbitals in the plane overlap with the p orbitals of the left-hand oxygen atom. 

The bonds along the molecular axis of carbon dioxide can be thought of as forming from combinations of hybrid orbitals on each of the three atoms. Assume the molecule lies on the z axis in the arrangement OaCOb. 

Depending on the used medium, the side products (detonation carbon) produced in detonation synthesis contain 20-70 wt% diamond. The remainder is a mixture of various structural forms of ip -carbon, the state of hybridized electronic orbitals characteristic of graphite. The yield of the diamond phase is the highest with carbon dioxide in the solid state, and the lowest, when it is in gaseous form. The medium most frequently used in industrial synthesis is water. 

For many years, the limited similarity between silicon and carbon excited the scientific community. Carbon and silicon share the same outer shell electronic structure, s, which permits sp hybridization and dominant tetrahedral coordination, as well as dominance of the tetravalent oxidation state. Nevertheless, silicon chemistry is markedly poorer compared to that of carbon. Double silicon bonds and silicon catenation are scarce, and crystalline silicon, which is so widely used in the electronics industry, is never encountered in nature. Instead, sUicon-oxygen bonds dominate natural silicon chanistry, and solid silica and silicates have no common physicochemical features with carbon dioxide and carbonates. The silicon atom is larger than carbon, it is less electronegative, has lower nuclear electric charge shielding and, perhaps most importantly, it has vacant d-orbitals in its outer shell all these dictate the reactivity of silicon. Several consequences of these differences are especially significant, and they are also relevant to sol-gel electrochemistry. 

Represent bonding in the carbon dioxide molecule, CO2, by (a) a Lewis structure and (b) the valence-bond method. Identify cr and it bonds, the necessary hybridization scheme, and orbital overlap. 

---