What is the average formal charge on oxygen

Pay close attention to the neutral forms of the elements below because that is how they will appear most of the time. Carbon, the most important element for organic chemists.

In the structures of methane, methanol, ethane, ethene, and ethyne, there are four bonds to the carbon atom. And each carbon atom has a formal charge of zero. In other words, carbon is tetravalent , meaning that it commonly forms four bonds.

Carbon is tetravalent in most organic molecules, but there are exceptions. Carbocations occur when a carbon has only three bonds and no lone pairs of electrons. Carbanions occur when the carbon atom has three bonds plus one lone pair of electrons. Two other possibilities are carbon radicals and carbenes, both of which have a formal charge of zero.

A carbon radical has three bonds and a single, unpaired electron. Carbon radicals have 4 valence electrons and a formal charge of zero. Carbenes are a highly reactive species, in which a carbon atom has two bonds and one lone pair of electrons, giving it a formal charge of zero. Though carbenes are rare, you will encounter them in section 8. You should certainly use the methods you have learned to check that these formal charges are correct for the examples given above.

More importantly, you will need, before you progress much further in your study of organic chemistry, to simply recognize these patterns and the patterns described below for other atoms and be able to identify carbons that bear positive and negative formal charges by a quick inspection. The common bonding pattern for hydrogen is easy: hydrogen atoms in organic molecules typically have only one bond, no unpaired electrons and a formal charge of zero.

The hydrogen radical is a hydrogen atom with no bonds, a single unpaired electron and a formal charge of 0. Nonetheless, the idea of a proton will be very important when we discuss acid-base chemistry, and the idea of a hydride ion will become very important much later in the book when we discuss organic oxidation and reduction reactions. As a rule, though, all hydrogen atoms in organic molecules have one bond, and no formal charge. The common arrangement of oxygen that has a formal charge of zero is when the oxygen atom has 2 bonds and 2 lone pairs.

All three patterns of oxygen fulfill the octet rule. If it has two bonds and two lone pairs, as in water, it will have a formal charge of zero. Oxygen can also exist as a radical, such as where an oxygen atom has one bond, two lone pairs, and one unpaired free radical electron, giving it a formal charge of zero. For now, however, concentrate on the three main non-radical examples, as these will account for most oxygen containing molecules you will encounter in organic chemistry.

Nitrogen has two major bonding patterns, both of which fulfill the octet rule:. If a nitrogen has three bonds and a lone pair, it has a formal charge of zero. In a fairly uncommon bonding pattern, negatively charged nitrogen has two bonds and two lone pairs. Two third row elements are commonly found in biological organic molecules: phosphorus and sulfur. Remember that elements in the third row of the periodic table have d orbitals in their valence shell as well as s and p orbitals, and thus are not bound by the octet rule.

The halogens fluorine, chlorine, bromine, and iodine are very important in laboratory and medicinal organic chemistry, but less common in naturally occurring organic molecules. Halogens in organic compounds usually are seen with one bond, three lone pairs, and a formal charge of zero.

Once you have gotten the hang of drawing Lewis structures, it is not always necessary to draw lone pairs on heteroatoms, as you can assume that the proper number of electrons are present around each atom to match the indicated formal charge or lack thereof. Occasionally, though, lone pairs are drawn if doing so helps to make an explanation more clear. The hydroxide ion, OH - , is drawn simply by showing the oxygen atom with its six valence electrons, then adding one more electron to account for the negative charge.

By changing the number of valence electrons the bonding characteristic of oxygen are now changed. Now the oxygen has three non-bonding lone pairs, and can only form one bond to a hydrogen. The oxygen has one non-bonding lone pair and three unpaired electrons which can be used to form bonds to three hydrogen atoms.

As an example of how formal charges can be used to determine the most stable Lewis structure for a substance, we can compare two possible structures for CO 2. Both structures conform to the rules for Lewis electron structures.

C is less electronegative than O, so it is the central atom. C has 4 valence electrons and each O has 6 valence electrons, for a total of 16 valence electrons. What is the formal charge on each atom in the tetrahydridoborate ion?

What is the relationship between having full valence shells and formal charges? Do full valence shells always result in a formal charge of zero? Why is formal charge used? See all questions in Formal Charge. Impact of this question views around the world. Using Formal Charge to Determine Molecular Structure Nitrous oxide, N 2 O, commonly known as laughing gas, is used as an anesthetic in minor surgeries, such as the routine extraction of wisdom teeth.

Which is the likely structure for nitrous oxide? Solution Determining formal charge yields the following:. The structure with a terminal oxygen atom best satisfies the criteria for the most stable distribution of formal charge:. The number of atoms with formal charges are minimized Guideline 2 , and there is no formal charge larger than one Guideline 2. This is again consistent with the preference for having the less electronegative atom in the central position.

You may have noticed that the nitrite anion in Example 3 can have two possible structures with the atoms in the same positions. The electrons involved in the N—O double bond, however, are in different positions:.

If nitrite ions do indeed contain a single and a double bond, we would expect for the two bond lengths to be different. A double bond between two atoms is shorter and stronger than a single bond between the same two atoms. Instead, we use the concept of resonance : if two or more Lewis structures with the same arrangement of atoms can be written for a molecule or ion, the actual distribution of electrons is an average of that shown by the various Lewis structures.

We call the individual Lewis structures resonance forms. The actual electronic structure of the molecule the average of the resonance forms is called a resonance hybrid of the individual resonance forms.

A double-headed arrow between Lewis structures indicates that they are resonance forms. We should remember that a molecule described as a resonance hybrid never possesses an electronic structure described by either resonance form. It does not fluctuate between resonance forms; rather, the actual electronic structure is always the average of that shown by all resonance forms. George Wheland, one of the pioneers of resonance theory, used a historical analogy to describe the relationship between resonance forms and resonance hybrids.

A medieval traveler, having never before seen a rhinoceros, described it as a hybrid of a dragon and a unicorn because it had many properties in common with both. Just as a rhinoceros is neither a dragon sometimes nor a unicorn at other times, a resonance hybrid is neither of its resonance forms at any given time. Like a rhinoceros, it is a real entity that experimental evidence has shown to exist. It has some characteristics in common with its resonance forms, but the resonance forms themselves are convenient, imaginary images like the unicorn and the dragon.

One oxygen atom must have a double bond to carbon to complete the octet on the central atom. All oxygen atoms, however, are equivalent, and the double bond could form from any one of the three atoms.

This gives rise to three resonance forms of the carbonate ion. Because we can write three identical resonance structures, we know that the actual arrangement of electrons in the carbonate ion is the average of the three structures. Again, experiments show that all three C—O bonds are exactly the same.

The online Lewis Structure Make includes many examples to practice drawing resonance structures. In a Lewis structure, formal charges can be assigned to each atom by treating each bond as if one-half of the electrons are assigned to each atom. These hypothetical formal charges are a guide to determining the most appropriate Lewis structure. A structure in which the formal charges are as close to zero as possible is preferred.

Resonance occurs in cases where two or more Lewis structures with identical arrangements of atoms but different distributions of electrons can be written.