Unit 2: Topic 6 - Resonance and Formal Charge Flashcards
Define Resonance.
Resonance is used to describe delocalized electrons which occurs when one Lewis structure is not enough to describe the molecule (or ion). A delocalized electron is simply an electron that does not directly associate with a covalent bond or a specific atom.
When depicting a molecule or ion requiring resonance structures, ALL structures have to be shown with double headed arrows between each. This is because the correct Lewis structure is actually the average of all possible resonance structures.
If a structure can be represented by multiple Lewis structures, explain the importance of depicting all possible Lewis structures (refinement)
If a structure has multiple possible Lewis structures, there is a chance that the true structure is impossible to draw using only single, double, or triple bonds. In fact, it is likely that the true structure is the average of all possible Lewis structures, and that bond lengths are between that of single and double bonds (hence the importance of refinement). For example, consider carbonate ion, CO3 2-. If we were to draw out the Lewis structure, we would notice that central carbon would have two single bonds to oxygen and one double bond to oxygen (in a trigonal planar shape). However, all three bond lengths are 136pm, which contradicts the idea of two single bonds and one double bond. However, if we take all three possible structures and average them out (so the true bond is something between a single bond and a double bond), then this is a refinement of a Lewis structure.
Consider ozone, O3. Draw out the resonance structures for ozone, estimate the percentage that each resonance structure makes up of the overall molecule, and qualitatively explain the bond structure.
Since all atoms are oxygen, we draw out three oxygen atoms. First, assume single bonds between the central oxygen with the other oxygen atoms: since there are two single bonds this way, 4 electrons have been used so far. To finish the octets on the terminal oxygens would require us to place 6 electrons (3 lone pairs) on each, using 12 electrons. O3 has 18 total valence electrons and we have used up 16, so the two remaining electrons go on the central oxygen. However, the central atom is not a full octet. This can be fixed by changing either one of the single bonds into a double bond as shown here. Since O3 is the average of these two resonance structures, the bond lengths are between that of a single and double bond.
Define formal charges and why they are important when choosing a Lewis structure for a molecule.
The formal charge of an atom is the net charge an atom would bear if all bonds to this atom were nonpolar covalent. To calculate this value, we can do (formal charge) = (number of valence electrons in free orbital) - (number of lone pair electrons) - 1/2(number of bond pair electrons). Formal charges are used to determine the best possible structure for a molecule or ion. The structure with more covalent bonds, the least number of formal charges (all close to 0), and least separation is more stable. If there is still multiple possibilities, the structure with negative charge on the most electronegative atom is the most stable.
Consider thiocyanate ion, SCN-. Use formal charges to decide which Lewis structure is optimal.
Diagram
Carbon is implied to be the central atom. Looking at the diagram, we can see that the middle choice is most stable, because the formal charges are closest to 0 and nitrogen, the most electronegative, has the only negative formal charge.
Consider nitrous oxide, N2O. Use formal charges to decide which Lewis structure is optimal.
It’s a bit harder to see what the central atom should be. However, we can consider the octet rule. Nitrous oxide has 16 valence electrons. If the central atom is single bonded to the other two atoms (i.e. two total bonds), there would be 8 lone pairs, three on each terminal atom, and two on the central atom. However, this is too many electrons. This can be reduced by giving the central atom 4 total bonds: now there are no lone pairs on the central atom. Therefore, the central atom should be nitrogen, as the formal charge would be +1, whereas with oxygen it would be +2. Now, since the central formal charge is positive and the overall charge is neutral, the remaining formal charges should be negative. We would like the terminal nitrogen to be +0 and the terminal oxygen to be -1 formal charges, as the more electronegative atom should have the negative formal charge (if possible). If we give the central nitrogen a triple bond to the terminal nitrogen, we will notice that the formal charge of terminal nitrogen is 5 - 2 - 1/2(6) = 0, which is desired. The final Lewis structure is shown here.
Nitrogen Monoxide (nitric oxide), NO, has an odd number of valence electrons. What would the optimal Lewis structure be?
The octet rule cannot fully apply here: there is an odd number of total electrons, so inevitably one will be unpaired. We start with a single bond between nitrogen and oxygen. Since oxygen is more electronegative, we include 3 lone pairs on oxygen first. After this, there are 3 more electrons to be placed on nitrogen. Nitrogen here has 5 electrons: although we have discovered that the octet rule will be violated, we want there to be as many electrons on nitrogen as possible without having more than 8. This is possible by changing the single bond to a double bond. link
Draw the Lewis Structure for NO2.
Note: This Lewis structure also requires resonance.
NO2 has 17 valence electrons. The less electronegative atom, nitrogen, should be the central atom. Consider first if we single bond the nitrogen to each terminal oxygen. Then, each terminal oxygen would have 3 lone pairs, leaving only 3 electrons for the central nitrogen. Nitrogen now only has 5 electrons. If we change one of the single bonds to a double bond, nitrogen will have 7, which is as close to the octet as possible. Notice that since either bond can be the single bond and the other will be the double bond, this structure will have two resonance structures.
