0.5 Covalent bonding and electron pair sharing  (Page 4/8)

It is important to note that there exist no hydrocarbons where the number of hydrogens exceeds two more thantwice the number of carbons. For example,

$C{H}_5$ does not exist, nor does $C_2{H}_8$ . We correspondingly find that all attempts to draw Lewis structureswhich are consistent with the octet rule will fail for these molecules. Similarly, $C{H}_3$ and $C_2{H}_5$ are observed to be so extremely reactive that it is impossible to prepare stable quantities of either compound. Again we find that itis not possible to draw Lewis structures for these molecules which obey the octet rule.

We conclude from these examples that, when it is possible to draw a Lewis structure in which each carbon has acomplete octet of electrons in its valence shell, the corresponding molecule will be stable and the hydrocarbon compound will existunder ordinary conditions. After working a few examples, it is apparent that this always holds for compounds with molecularformula $C_n{H}_{2n+2}$ .

On the other hand, there are many stable hydrocarbon compounds with molecular formulae which do not fit theform $C_n{H}_{2n+2}$ , particularly where the number of hydrogens is less than $2n+2$ . In these compounds, the valences of the carbon atoms are not quiteso obviously satisfied by electron pair sharing. For example, in ethene $C_2{H}_4$ and acetylene $C_2{H}_2$ there are not enough hydrogen atoms to permit each carbon atom to be bonded to four atoms each. In each molecule, the two carbonatoms must be bonded to one another. By simply arranging the electrons so that the carbon atoms share a single pair ofelectrons, we wind up with rather unsatisfying Lewis structures for ethene and acetylene, shown here .

Note that, in these structures, neither carbon atom has a complete octet of valence shell electrons. Moreover,these structures indicate that the carbon-carbon bonds in ethane, ethene, and acetylene should be very similar, since in each case asingle pair of electrons is shared by the two carbons. However, these bonds are observed to be chemically and physically verydifferent. First, we can compare the energy required to break each bond (the bond energy or bond strength ). We find that the carbon-carbon bond energy is 347 kJ in $C_2{H}_6$ , 589 kJ in $C_2{H}_4$ , and 962 kJ in $C_2{H}_2$ . Second, it is possible to observe the distance between the twocarbon atoms, which is referred to as the bond length . It is found that carbon-carbon bond length is 154 pm in $C_2{H}_6$ , 134 pm in $C_2{H}_4$ , and 120 pm in $C_2{H}_2$ . ( $1\mathrm{picometer}=1\mathrm{pm}=10^{-12}m$ ). These observations reveal clearly that the bonding between thecarbon atoms in these three molecules must be very different.

Note that the bond in ethene is about one and a half times as strong as the bond in ethane; this suggests thatthe two unpaired and unshared electrons in the ethene structure above are also paired and shared as a second bond between the twocarbon atoms. Similarly, since the bond in acetylene is about two and a half times stronger than the bond in ethane, we can imaginethat this results from the sharing of three pairs of electrons between the two carbon atoms. These assumptions produce the Lewisstructures here .