Carbon formal charges

Isomtriles are stable often naturally occumng compounds that contain a divalent carbon An example is axisonitnle 3 which can be isolated from a species of sponge and possesses anti malanal activity Write a resonance form for axisonitnle 3 that satisfies the octet rule Don t for get to include formal charges... 

There are several ways to choose the more plausible of two structures differing in their arrangement of atoms. As pointed out in Example 7.1, the fact that carbon almost always forms four bonds leads to the correct structure for ethane. Another approach involves a concept called formal charge, which can be applied to any atom within a Lewis structure. The formal charge is the difference between the number of valence electrons in the free atom and the number assigned to that atom in the Lewis structure. The assigned electrons include—... 

To show how this works, let s calculate the formal charges of carbon and oxygen in the two structures written above for methyl alcohol ... 

C bond with two oxygen atoms attached to each carbon atom) (b) BrO4 (c) the acetylide ion, C22. Assign formal charges to each atom. 

A formal charge is a charge associated with an atom that does not exhibit the expected number of valence electrons. When calculating the formal charge on an atom, we first need to know the number of valence electrons the atom is supposed to have. We can get this number by inspecting the periodic table, since each column of the periodic table indicates the number of expected valence electrons (valence electrons are the electrons in the valence shell, or the outermost shell of electrons— you probably remember this from high school chemistry). For example, carbon is in Column 4A, and therefore has four valence electrons. Now you know how to determine how many electrons the atom is supposed to have. 

Now we are in a position to compare how many valence electrons the atom is supposed to have (in this case, four) with how many valence electrons it actually has (in this case, four). Since these numbers are the same, the carbon atom has no formal charge. This will be the case for most of the atoms in the structures you will draw in this course. But in some cases, there will be a difference between the number of electrons the atom is supposed to have and the number of electrons the atom actually has. In those cases, there will be a formal charge. So let s see an example of an atom that has a formal charge. 

If carbon bears a formal charge, then we cannot just assume the carbon has four bonds. In fact, it will have only three. Let s see why. Let s first consider C+, and then we will move on to C . 

If carbon has a positive formal charge, then it has only three electrons (it is supposed to have four electrons, because carbon is in Column 4A of the periodic table). Since it has only three electrons, it can form only three bonds. That s it. So, a carbon with a positive formal charge will have only three bonds, and you should keep this in mind when counting hydrogen atoms ... 

Now let s consider what happens when we have a carbon atom with a negative formal charge. The reason it has a negative formal charge is because it has one more electron than it is supposed to have. Therefore, it has live electrons. Two of these electrons form a lone pair, and the other three electrons are used to form bonds ... 

From all of the cases above (oxygen, nitrogen, carbon), you can see why you have to know how many lone pairs there are on an atom in order to figure out the formal charge on that atom. Similarly, you have to know the formal charge to figure out how many lone pairs there are on an atom. Take the case below with the nitrogen atom shown ... 

Answer We read the arrows to see what is happening. One of the lone pairs on oxygen is coming down to form a bond, and the C=C double bond is being pushed to form a lone pair on a carbon atom. This is very similar to the example we just saw. We just get rid of one lone pair on oxygen, place a double bond between carbon and oxygen, get rid of the carbon-carbon double bond, and place a lone pair on carbon. Finally, we must draw any formal charges ... 

Look carefully at the formal charges. The negative charge used to be on oxygen, but now it moved to carbon. 

When we treat all bonds as covalent, the carbon atom appears to have four electrons of its own. Carbon is supposed to have four valence electrons. When we compare how many electrons carbon actually has with the number of electrons it is supposed to have, we see that everything is just right in this case. It is supposed to have four valence electrons, and it is clearly using four valence electrons. Therefore, there is no formal charge. 

Each carbon atom has four bonds, and each oxygen atom has zero formal charge. The — CO2H group, with one CDO double bond and an acidic C—O—H linkage, is characteristic of carboxylic acids. 

Notice that the zinc atom is associated with only four valence electrons. Although this is less than an octet, the adjacent carbon atoms have no lone pairs available to form multiple bonds. In addition, the formal charge on the zinc atom is zero. Thus, Zn has only four electrons in the optimal Lewis structure of dimethyizinc. This Lewis stmcture shows two pairs of bonding electrons and no lone pairs on the inner atom, so Zn has a steric number of 2. Two pairs of electrons are kept farthest apart when they are arranged along a line. Thus, the C—Zn—C bond angle is 180°, and linear geometry exists around the zinc atom. 

C09-0107. Write Lewis structures and calculate formal charges for the following polyatomic ions (a) bromate (b) nitrite (c) phosphate and (d) hydrogen carbonate. 

C09-0108. Carbon, nitrogen, and oxygen form two different polyatomic ions cyanate ion, NCO, and isocyanate ion, CNO". Write Lewis stmctures for each anion, including near-equivalent resonance structures and indicating formal charges. 

Both structures II and III have an arrangement of atoms that places a positive formal charge on atoms that are higher in electronegativity than carbon. Consequently, the most stable arrangement of atoms is as shown in structure I. Some compounds containing the ion having structure III (the fulminate ion) are known, but they are much less stable than the cyanates (structure I). In fact, mercury fulminate has been used as a detonator. 

The BH3 molecule is not stable as a separate entity. This molecule can be stabilized by combining it with another molecule that can donate a pair of electrons (indicated as ) to the boron atom to complete the octet (see Chapter 9). For example, the reaction between pyridine and B2H6 produces C5H5N BH3. Another stable adduct is carbonyl borane, OC BH3 in which a pair of electrons is donated from carbon monoxide, which stabilizes borane. In CO, the carbon atom has a negative formal charge, so it is the "electron-rich" end of the molecule. Because the stable compound is B2H6 rather than BH3, the bonding in that molecule should be explained. 

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