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Lewis Structures
Lewis Structures
In the previous section, we used the valences of atoms to predict the formulas and structures they would adopt when bonding. This approach relies on two assumptions. First, that the octet rule is always obeyed. Second, that bonding always involves one electron from atom A and one electron from atom B.
These assumptions are, however, not always valid. The structure of SO3 shown at right is a good example of how they are both broken.
First, looking at the sulfur, you can see that it has blown way past the octet rule. With six bonding pairs of electrons, it has twelve valence electrons. This is a thing sulfur and other elements do sometimes.
Second, if we just lined up S and O atoms and started pairing their electrons up, this is not necessarily the structure we would expect to get. In this section, we will see that it is often more useful to think of electrons as belonging to a whole molecule than to a single atom.
Steps For Drawing Lewis Structures
Steps For Drawing Lewis Structures
Here are some guidelines for taking a formula of a covalently bonded molecule and drawing its Lewis dot structure.
*Note: These steps assume molecules with a simple structure, meaning one atom in the middle and the rest connected only to that atom. You will not be asked to draw anything with a more complicated structure, though playing around with such examples may be helpful in honing your skills. Ask your instructor for challenge problems like these, if you want.
- Identify Central Atom
Say you are given the formula NBr3 and asked to draw its Lewis structure. The first thing to ask yourself is "what atom goes in the middle of my molecule?"
There are a few clues you can use to establish this. Often, the central atom will be listed first in the formula. Here we can see that the nitrogen is listed first. A better clue, though, is position in the periodic table. In general, atoms further to the left in the table tend to be in the center. If two atoms are in the same column, the one further down is more likely to be in the center. In this case, we conclude nitrogen to be the central atom.
2. Count Valence Electrons
2. Count Valence Electrons
Each atom in a molecule brings valence electrons with it. These valence electrons should appear in your final structure; no more, no less. In this case, we have one nitrogen with five valence electrons, and three bromines with seven valence electrons each. 5 + 3*7 = 26 valence electrons
3. Draw Single Bonds
3. Draw Single Bonds
For starters, our atoms must all be connected, so we begin by drawing single bonds from the outer atoms to the central atom.
We should keep track of our valence electrons as they are placed. These three bonds use up six valence electrons, so we have 20 valence electrons remaining.
4. Fill Octets on Outer Atoms
4. Fill Octets on Outer Atoms
We will need to eventually give each atom eight valence electrons, consistent with the octet rule. We tend first to our outer atoms. Each bromine atom has two electrons from its bond, so we give them all six more to get to eight. This uses 18 valence electrons, leaving 2 valence electrons.
*Note: If hydrogen is involved, it wants a duet, not an octet, so it gets left alone at this step
5. Place Leftover Electrons on Central Atom
5. Place Leftover Electrons on Central Atom
This step will not always be necessary, as sometimes you will use all your electrons up in Step 4. However, if there are any left, at this point they should go on the central atom. We have 2 remaining, so we put them on the nitrogen.
Examining our structure, we can see we have eight valence electrons on all our atoms. So our molecule is finished!
6. Make Multiple Bonds if Needed
6. Make Multiple Bonds if Needed
In this case, we did not need any double or triple bonds. However, some molecules will require them. Here's the example of HCN, which I have gotten to the end of Step 5. We have found our central atom, placed single bonds, and filled the octet on nitrogen (and the duet on hydrogen).
This has used up all ten of our valence electrons, and carbon is still without a full octet; it has only four valence electrons. The way we allow it to get an octet is by taking lone pairs from nitrogen and turning them into multiple bonds. In this case, a triple bond is needed.
It's very important that you not just start drawing double and triple bonds willy-nilly. They should only be drawn if the central atom is short of a full octet. Furthermore, they can only be made by taking lone pairs from outer atoms and moving them between the atoms to make a bond.
Practice
Practice
You should now practice using this procedure to draw Lewis structures for molecules. Here are some example formulas to get started on. You can scroll down to see the answers and check your work.
SCl2
N2O
SeO2
O3
Answers
Answers
Octet Rule Violators
Octet Rule Violators
I mentioned previously that the octet rule is not as hard and fast as it often seems. I showed you the example of SO3 above. Here I will tell you about other octet rule violations you need to be aware of.
Hypovalence
Hypovalence
Not all elements that covalently bond are trying to get to a state of eight valence electrons. For one element in particular, boron, six valence electrons seems to be a perfectly acceptable number.
Thus, we see compounds like the one at right, in which boron only makes three single bonds. This type of situation, with fewer than eight valence electrons, is called hypovalence ("hypo" meaning "beneath").
Hypervalence
Hypervalence
The opposite of hypovalence is hypervalence - where an atom has more than eight valence electrons.
Hypervalence is never observed for second-row elements like C, N, O, or F. However, it is fairly common in elements of the third row and above, meaning P, S, Cl, Se, Br, and I. All of these elements form some compound or another in which they have ten or even twelve valence electrons. An example, phosphorus with five bonds to fluorine, is shown at right.
Radicals
Radicals
In almost all the compounds we have seen so far, almost all the electrons have come in pairs - either bonding pairs or lone pairs. This is by far the norm in the chemistry of covalent bonding. However, there are rare exceptions. These are known as radicals; you will not need to do much of anything in terms of drawing their structures, but I want you to be aware of their existence as you learn about the diversity of forms chemical bonding can product. An example is shown at right.
Polyatomic Ions
Polyatomic Ions
We have discussed ionic bonding in Lesson 7. In that lesson, you learned that there are a number of structures called polyatomic ions, which are clusters of atoms that carry a charge.
These clusters are held together by covalent bonding; in fact, they are very much like molecules, just with excess or deficient electrons.
Lewis structures can be drawn for polyatomic ions. At right is the Lewis structure for carbonate(CO32−). If drawing a polyatomic ion from its formula, you follow the same instructions given above. The difference is, you add or subtract electrons based on the charge.
Carbon has four valence electrons, oxygen has six each (times three), and the charge on carbonate is -2, meaning it has two excess electrons, for a total of 4 + 3*6 + 2 = 24. Count them up in the image at right and confirm they are all there.
So, in an ionic compound with polyatomic ions in it, there is both ionic and covalent bonding occurring. For example, in Na2CO3, the polyatomic ions are held together internally by covalent bonds, while those ions are bound to the sodium ions by ionic bonds.
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