D7.2 When One “Bond” is not Enough
In space-filling models like those of ethene (C2H4) and ethyne (C2H2):
atoms appear connected, but there’s no visible indication of how many electron pairs are shared. These 3D models communicate shape and size, but not the details of electron arrangement.
That’s where Lewis structures come in. Lewis structures are built from electron bookkeeping: we assign valence electrons to atoms and connect them in ways that reflect their usual bonding patterns. Sometimes, after placing all single bonds, we find that all the valence electrons have not been used. When this happens, Lewis structures introduce multiple bonds, two or three shared electron pairs between atoms, to ensure all valence atoms are accounted for.
For example:
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Activity
Draw the valid Lewis structures for ethane (C2H6), ethene (C2H4), and ethyne (C2H2). What patterns do you notice in how many pairs of electrons are shared between the carbon atoms?
Draw and Write in your notebook, then left-click here for an explanation.
All three molecules have two carbon atoms bonded to each other. As the number of hydrogen atoms that are bonded to each carbon atom decrease, the number of electron pairs shared between the carbon atoms increase.
Even though single, double, and triple bonds are conventions used in Lewis structures, they reflect real differences in molecular behavior:
- As more electrons are shared, bond strength increases and bond length decreases.
Average bond strength Average bond length C–C 346 kJ/mol 154 pm C=C 598 kJ/mol 134 pm C≡C 813 kJ/mol 121 pm - Double and triple bonds often restrict rotation and influence reactivity.
- They affect electron distribution, which influences how molecules interact.
So while molecules just have bonds (electron-sharing interactions between atoms), our descriptions of those bonds (single, double, triple) are powerful tools for reasoning about structure, stability, and reactivity.
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