Chapter Electrochemistry. Chapter Radioactivity and Nuclear Chemistry. Chapter Transition Metals and Coordination Complexes. Chapter Biochemistry. Full Table of Contents. This is a sample clip. Sign in or start your free trial. JoVE Core Chemistry. Previous Video Next Video. Next Video 9. Embed Share. Some molecules or polyatomic ions can be represented by multiple Lewis structures, but how to decide which one is the dominant structure?
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Save to playlist. So let's start with this hydrogen over here. So what's the number of valence electrons in a free, neutral atom of hydrogen? Well we've seen this multiple times, you could look at this on the periodic table of elements, free neutral hydrogen has one valence electron. Now how many valence electrons are allocated to the bonded atom? Well one way to think about it is, draw a circle around that atom in the molecule, and you want to capture all of the lone pairs, and you want to capture, you can think of it as half the bond, you could say for each bond, it's going to be one electron 'cause it's half of the shared electrons, each bond is two shared electrons, but you're gonna say half of those, and then you have no lone pairs over here, so the number of valence electrons allocated to bonded atom, in the case of hydrogen here, is one, and so we are dealing with a formal charge of zero for this hydrogen.
Now what about this oxygen here? Well we do the same exercise, I like to draw a little bit of a circle around it. And so the number of valence electrons in a free, neutral oxygen we've seen multiple times, that is six, and then from that, we're going to subtract the number of valence electrons allocated to the bonded atom. So the bonded atom has two lone pair electrons, and then it gets half of the shared electrons, so half of the shared electrons would be one from this bond, one from that bond, and one from that bond.
So you add them all together, two, three, four, five. So six minus five is equal to positive one, and so the formal charge on this oxygen atom, in this configuration of nitrous acid is positive one. 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. Dividing the remaining electrons between the O atoms gives three lone pairs on each atom:. This structure has an octet of electrons around each O atom but only 4 electrons around the C atom.
No electrons are left for the central atom. To give the carbon atom an octet of electrons, we can convert two of the lone pairs on the oxygen atoms to bonding electron pairs. There are, however, two ways to do this. We can either take one electron pair from each oxygen to form a symmetrical structure or take both electron pairs from a single oxygen atom to give an asymmetrical structure:. Both Lewis electron structures give all three atoms an octet. How do we decide between these two possibilities?
The formal charges for the two Lewis electron structures of CO 2 are as follows:. Thus the symmetrical Lewis structure on the left is predicted to be more stable, and it is, in fact, the structure observed experimentally. Remember, though, that formal charges do not represent the actual charges on atoms in a molecule or ion.
They are used simply as a bookkeeping method for predicting the most stable Lewis structure for a compound. The Lewis structure with the set of formal charges closest to zero is usually the most stable. Draw two possible structures, assign formal charges on all atoms in both, and decide which is the preferred arrangement of electrons. Asked for: Lewis electron structures, formal charges, and preferred arrangement. B Calculate the formal charge on each atom using Equation 2. Formal charge is the actual charge on an individual atom within a larger molecule or polyatomic ion.
The sum of formal charges on any molecule or ion results in the net overall charge. This concept is simple enough for small ions. Chloride obviously has a negative charge. Even the negative charge on the hydroxide oxygen is simple to understand.
But what if you have a much larger group of bound atoms with an overall net charge? For example, the negative nitrate or triple negative phosphate. Think back to general chemistry when you studied ion formation. An ion is simply an atom or molecule that gained or lost electrons to get a net charge. If the atom has just one more negative electron, than protons, it will have a net negative charge.
When I first studied formal charge I was lost. The formula in my textbook was long, tedious and brutal. Do you dread running through this every time you finish a reaction?
There are just too many steps and calculations. Here another option, MY version — one that is easier, faster, and comes out with the same result! Definitely faster, right? Each bond only counts for a single electron since the second electron in the bond is touching the other atom. Want to see this shortcut brought to life? See my formal charge video below.
When you first study formal charge it helps to draw out the Lewis Structure for every molecule in question. As you work through Lewis Structures it helps to have a periodic table handy.
Since we are working on saving you time though, your most efficient results will involve you memorizing the following 10 atoms. As you memorize their order and location, pay particular attention to these trends:. Following the checklist in the Lewis Structures video we have oxygen bound to hydrogen, with 3 lone pairs around oxygen. Following the checklist we draw our atoms, bonds, and electrons. Due to resonance we would show three structures for nitrate.
Neutral nitrogen should have 5 valence electrons but our drawing only shows 4 attached. The double bound oxygen is happy, stable, and has a net neutral charge. Finally we have 2 single bound oxygen atoms with 3 lone pairs each.
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