Too Many Electrons. Lone Pairs

What happens to electrons which are left over after all bonds have been formed Do they associate with individual atoms or are they spread uniformly throughout the molecule Draw a Lewis structure for trimethylamine. How many electrons are needed to make bonds How many are left over Where are they Display the highest-occupied molecular orbital (HOMO) for trimethylamine. Where is it located  

Examine the HOMO-2 forphenylisocyanide. Is it directly involved in any o or K bonds If so, which bonds If not, describe where it is located. Draw a Lewis structure for the molecule which is consistent with your result. 

Electrons which are left over after bonds have been formed are referred to as lone pairs . Although they may not contribute (directly) to bonding, they do take up space. You camiot actually see lone pairs, but you can see the space which they occupy and infer whether or not they contribute significantly to a molecule s overall size and shape. 

Draw Lewis structures for methyl anion, ammonia and hydronium cation. How many electrons are left over in each after all bonds have been made Display and compare electron density surfaces for methyl anion, ammonia and hydronium cation. Which is the smallest molecule Which is the largest Rationalize your observation. (Hint Compare the number of electrons in each molecule, and the nuclear charge on the central atom in each molecule.) 

HOMO for trimethylamine shows location of electrons which are left over after all bonds have been formed. 

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Step 4 Calculate the number of valence electrons in your tentative diagram and compare it with the actual number of valence electrons. If the tentative diagram has too many electrons, remove a lone pair from the central atom and from a terminal atom, and replace them with an additional bonding pair between those atoms. If the tentative diagram still has too many electrons, repeat the process. 

This time there are too many electrons in the tentative diagram. Modify the tentative diagram by removing two lone pairs and replacing them with one bonding pair. 

There are 4 too many electrons in the tentative diagram this time however, the remedy is the same. Erase two lone pairs from the carbons and replace them with a bonding pair to reduce the electron count by 2. Then do it again to reduce the electron count by a total of 4. The result is a triple bond. 

It should be clear from the above that because of sustained efforts over many decades, significant progress has now been achieved in the understanding of the freezing of water into ice. However, there stiU remain many unsolved problems in this area. For example, we do not yet have a quantitative theory of the nucleation of ice in supercooled water. The molecular models we use in simulations are perhaps too primitive, as most of them do not include the polarizabihty of water molecules. The polarizability of water is large due to the two lone pairs of electrons on the lone oxygen atom. Perhaps one would need to consider quantum simulations to fully understand the freezing of ice. 

The number of valence electrons in the valence shells of early TMs (e.g.. Sc, Ti, V) is small, and hence, the metals would need too many ligands to acquire 18e. But then the sphere of the TM would be too crowded and the electrons of the various TM—L bonds would repel one another too strongly to afford these many ligands. Therefore, these TMs take fewer ligands and form complexes with fewer than 18e in the valence shell of the TM. As we move toward the end of the TM period (e.g., Ni, Cu, Pd, Ag, Pt, Au), the number of electrons on the TM becomes quite large, and those that remain as lone pairs on the TM repel the bonding electrons and weaken thereby the propensity of the TM to acquire 18e. 

In the following case an arrow is used to depict a potential resonance contributing structure of nitromethane. However, the result is a nitrogen atom with 10 electrons in its valence shell because there are too many bonds to N. Such mistakes can be avoided by remembering to draw all bonds and lone pairs on an atom so that the total number of electrons in each atom s valence shell is apparent. 

The structure above has six electrons in covalent bonds and eight electrons in four lone pairs, for a total of 14 electrons. The structure has two valence electrons too many. 

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