Which Electrons Do Lewis Structure Show?

which electrons do lewis structure show?

Lewis Structures for O-chem « reactions & distractions

Being able to draw Lewis structures is an important skill in organic chemistry. Being able to draw Lewis structures quickly and easily is invaluable. In this article, I will outline a few different methods for drawing Lewis structures, and you can decide for yourself which is easiest.

If you are taking an undergraduate organic chemistry class, I'm assuming you've already had some experience with drawing Lewis dot diagrams from Gen. Chem. Therefore, I'm only going outline the basic steps for each method but won't go into any great detail beyond that.

Method #1: The "dot" method (aka, the "jigsaw puzzle" method)

In this method, you place dots (representing the valence electrons) around each atom in the compound, and then piece them together (like a jigsaw puzzle) until all atoms have their octet rule satisfied (2e- around any Hydrogens, 8e- around most other 2nd-row elements, except Boron. Boron is weird.)

1. Draw all the atoms in the compound with dots representing valence electrons:

NOTE: If you are dealing with ions, rather than neutral compounds, you must remember to add in or subtract out the appropriate number of electron "dots" in order to make this work.

2. Put the atoms together by electron pairs. That is, arrange the atoms with their dots such that all the atoms are joined together with two dots in between (one from each atom). When you pair up two electrons from two different atoms, those electrons will form a bond, which can be re-drawn as a line. Electrons on atoms which are not part of bonds remain as dots (as with N, O, and F in the chart below.)

3. What happens when you have extra, unpaired electrons on some atoms? See if you can use those to form double- or triple-bonds! Below are three examples of compounds which have two carbon atoms and various numbers of hydrogen atoms. The two carbons are colored differently to help keep track of which electrons come from which carbon. In the first example, Ethane, there are just enough hydrogen atoms to completely bond with all the electrons provided by the two carbon atoms. In the second example, Ethene, after joining the two carbons together and adding in the hydrogens, there are two solitary electrons left over, one on each carbon. These two solitary electrons are joined to each other to form a double-bond between the carbons. A similar thing happens with the third example, Ethyne, which has enough solitary electrons left over to form a triple-bond between the carbons.

4. Hopefully you are done by this point. If not — tough luck. Keep trying different arrangements until you find one that works, or use method #2 below to help you.

Some other sites that may help:

Unfortunately, this method is probably the worst (IMHO) way to try to draw a Lewis structure for anything but the most elementary compounds (such as the above examples). For most organic compounds, the third method below is the one that works best for me.

Method #2: The "math" method:

As there is already a very good, step-by-step procedure for this method, I'm only going to go over the basics here.

1. Determine the total number of valence electrons in the molecule:

CH₄O
C has 4, times 1 = 4
H has 1, times 4 = 4
O has 6, times 1 = 6
total electrons = 14

* for ions, add or subtract the appropriate number of electrons from this total.

2. Determine the total number of electrons the molecule wants to have a "full" valence. (Remember: H wants 2, most 2nd-row elements want 8. Boron wants 6 'cause it's weird.)

CH₄O
C wants 8, times 1 = 8
H wants 2, times 4 = 8
O wants 8, times 1 = 8
full valence = 24

3. Subtract Step 2 from Step 1 to find how many electrons need to be shared between atoms to try to achieve a full valence.

24 wanted for full valence – 14 actual electrons = 10 need to be shared.

4. Remember that shared electrons form bonds, and each bond is made of two electrons. Divide the number of shared electrons from Step 3 by 2, to find the number of bonds.

10 electrons to be shared / 2 = 5 bonds total

5. Draw a structure for the molecule with the appropriate number of bonds between atoms. Place the least electronegative atoms in the center, as they will tend to need to form the most bonds. You may need to "play around" with the positions a bit until you find one that works.

CH₄O
H | H-C-O-H | H
is the best possible structure for this compound. (Can you name it?)

6. Note from Step 3 that we didn't use all of our electrons to form bonds. Note also, that some atoms in our structure from Step 5 don't yet have a full valence. This is where non-bonding "lone pair" electrons come into play: they are simply the electrons that are "left over" after the bonds are formed. Determine the number of non-bonding electrons by subtracting the number of shared electrons (Step 3) from the original total number of electrons (Step 1).

14 total electrons – 10 shared electrons = 4 non-bonding electrons

7. Non-bonding electrons also form pairs, so place them (in pairs) around any atoms which don't yet have a full valence in our structure from Step 5:

CH₄O
H | " H-C-O-H | " H
The 4 non-bonding electrons have been placed around the O, giving it a full valence.

8. Assign formal charges, if any. Formal charge on each atom is equal to its "normal" number of valence electrons, minus the number of electrons it "owns" in the molecule. An atom is considered to "own" half of its shared (bonding) electrons and all of its non-bonding electrons; in other words, count 1 for each bond and 2 for each lone pair.

CH₄O
C: 4 (valence) – 8/2 (shared) – 0 (non-bonding) = 0 fc
H: 1 (valence) – 2/2 (shared) – 0 (non-bonding) = 0 fc on each H
O: 6 (valence) – 4/2 (shared) – 4 (non-bonding) = 0 fc

(So I gave an easy example here…)

I like this method because it's thorough. However, it can also be quite time-consuming, especially when you get to Step 8 and find that you've made a mistake along the way somewhere and your formal charges don't add up, or that your formal charge assignments indicate that you haven't drawn the best possible structure for your compound. It's also pretty tough to do the calculations in your head, which you would need to do each time you draw a structure. (Imagine having to do that on a test!) Method #3 below takes all of the information we've learned from Method #1 and Method #2 and condenses it into a pattern that is fairly easy to memorize and utilize.

Method #3: The "pattern" method.

Okay, so I'm not normally an advocate of straight-up memorization in organic chemistry; however, I'll make an exception for this one. Learn the bonding patterns of the handful of elements most commonly encountered in o-chem (H, C, N, O, halogen, and sometimes B), and you'll be able to draw Lewis structures very quickly for just about any organic compound. The table below shows the bonds, lone pairs, and resulting geometry for Hydrogen as well as the five 2nd-row elements (B thru F) commonly found in organic chemistry.

Knowing the bonding pattern made by a particular element with a particular formal charge (fc) is simply a must when drawing organic compounds, particularly resonance structures. Learning these bonding patterns is actually easier than straight-up memorization, because there's a pattern in the above table — at least for the 2nd-row elements (Boron thru Fluorine). See if you can see it.

The pattern is that each neutral (0 formal charge) element has the same bonding pattern as the +1 formal charge species of the element to the right of itself on the periodic table (that is, the element with one higher atomic number), and also the same bonding pattern as the -1 formal charge species of the element to the left of itself on the periodic table (the element with one lower atomic number). For instance, Carbon with 0fc has the same bonding pattern +1fc Nitrogen and -1fc Boron:

Another thing to note about the bonding patterns depicted in the above table: they all show the bonding pattern for the maximum number of single bonds for that element/fc combo. However, it's fairly easy to account for double- or triple- bonds: simply count double-bonds as 2 bonds used up and triple-bonds as 3 bonds used up. This is illustrated below using Carbon with 0fc. Carbon normally will make 4 bonds and have no lone electron pairs. These 4 bonds can be 4 single-bonds to 4 different things, or perhaps a double-bond to one thing and 2 single-bonds to 2 other things, or maybe even a triple-bond to something and a single bond to something else.

(The other elements which have more than one bond shown in the table above behave similarly.)

That's all for now. Next time, we'll dig a bit deeper into how to use this method to draw Lewis structures.

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2012年1月29日日曜日

Which Electrons Do Lewis Structure Show?