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Unformatted text preview: ```````` Lewis Dot Structures
Prepared By: Richard E. Bleil, Ph.D. Introduction:
What can I say? There is more to life than receiving, and more than giving.
There is sharing. Yep, I'm talking about, of course, covalent compounds.
In ionic bonds, one element wants the electron so much more that it literally rips
that electron from the valence shell of the other. It's not quite so mean as it might sound,
because truth be told, the other atom is happy to lose that electron. This is an ionic bond,
wherein there is a complete transference of at least one electron from one atom (to
become the cation) to another (to become the anion). But, what happens when there is no
great difference between how much the two electrons want that electron? Each of them
want an electron, but neither is strong enough to just take it. When this happens, the
atoms are forced to share electrons.
Sharing electrons forms what we call a "covalent" bond. Just like in ionic
bonding, atoms share electrons to complete their valence shell.
The Octet Rule:
Long before quantum theory, it was noted that most elements will lose or share
electrons to get eight electrons in their valence shell. Now we know this is so as they are
trying to fill their "s" and "p" orbitals, but then it was called the "octet rule". Lewis Dot
structures are steeped in the "octet rule", which simply states that atoms will share
electrons to gain eight electrons in their valence shell. An exception to this readily
recognized is hydrogen. Since we now know that hydrogen's entire electronic
configuration is confined to the 1s orbital, it is not surprising that hydrogen will want
only two electrons, rather than eight. Early on, it was an empirical observation.
Another early empirical observation was that electrons preferred to exist in pairs.
I imagine it is the observation that atoms with unpaired electrons, called "radicals", are so
reactive. This observation led to the way in which we draw the Lewis dot structure for an
atom. We begin by writing the atomic symbol of the atom, and count the number of "s"
and "p" orbital valence shell electrons (eight maximum). We then draw one dot for each
of the s and p valence shell electrons in order in each of the four quadrants around the
symbol, top, right, bottom, left. Do NOT pair up electrons until we have one electron in
each of these quadrants. We then continue around the circle pairing up electrons until all
of these valence shell electrons are accounted for. 1 ```````` Each of the four quadrants, then, represents one of the four orbitals, the one s or three p
Now it becomes a game, a puzzle if you will, to match up the electrons so that
there are no unmatched electrons. The electrons with a partner are already happy, so we
need worry about the unpaired electrons only. Any bond must contain two electrons, but
between two atoms we can only form either single, double, or triple bonds. Since each
bond must contain two electrons, a single bond involves two electrons. Each of the two
bonds in a double bond involve two electrons each for a total of four electrons in a double
bond, and, similarly, there are six electrons involved in a triple bond, two for each of the
three bonds. Furthermore, although you can move around electrons (so that the unpaired
electrons or bonds are in different quadrants), it is important to remember that you must
NEVER add or remove electrons. It's a good idea to count the total number of valence
shell electrons at the beginning, and be sure they are all accounted for in the end with no
As usual, there are a few hints to keep in mind when working out Lewis dot
structures, just little things that seems to make them come together a little more easily.
These rules do not always apply, but they generally save a little time.
1. Start with the atom that needs the most electrons.
2. Start with single bonds; NEVER start a problem with multiple bonds.
3. If it is at all possible, avoid bonding one element to another element of the
same type. AN EXCEPTION to this is Carbon, which is actually quite happy
bonding with itself, so typically bond carbons together first. However, carbon
is the only atom that is really stable bonded to itself, so do not bond any other
atoms to themselves unless it is really necessary.
4. Save any atoms that only need one electron for the last step. This typically
includes hydrogen and any halogens.
5. Once all atoms are on the board except for those that need only one electron
are "on the board", count the number of unpaired electrons in the primary
structure. If there are the same number of electrons on the main structure as
elements remaining, start placing the rest of the elements. If there are more
unpaired electrons than remaining elements, start using multiple bonds, one
bond at a time, until the number of unpaired electrons is equal to the number
of remaining atoms.
6. Avoid making rings in your structures in general, but if there is a ring in the
structure, do now draw a chemical ring invloving fewer than five atoms.
7. Remember: you can ALWAYS start over with a new central arrangement if
What do you think? Should we start with a few examples? Let's start with a simple one.
Let's determine the Lewis-Dot structure for Hydrogen Astatine. 2 ```````` OK, so it's a pretty simplistic example, but let's see what we learned from it. First of all,
we notice that there are eight valence shell electrons total between both elements
involved initially (one from H, seven from the s and p orbitals of At), and we have eight
electrons in the final product; no more, no less. Notice that I used a dash to represent the
bond between H and At. A dash represents a pair of electrons just as a dot represents a
single electron. Is the octet rule satisfied in our final product? Notice that there are now
two electrons near the hydrogen (the only exception, remember?), and eight around
astatine. Notice that the shared electrons count towards both the hydrogen and the
astatine when checking the octet rule.
What does this structure tell us? Well, we see that there is a single bond between
the hydrogen and astatine, and we see that there are three lone pairs of electrons around
the astatine, none around the hydrogen. If YOUR answer looks different, it does not
matter provided you could tell me exactly what I just said. We have some freedom as to
how to draw these things, and all they tell us are bonding and lone pair information.
Remember: Lewis Dot Structures do NOT tell us any information on structure!
How about a more involved example? What is the Lewis-Dot structure of
CH3Cl? 3 ```````` We could have placed the chlorine in any of the four quadrants; remember, Lewis dot
structures do not speak to the actual shape of the molecule. However, we see here that
we have three hydrogens and one chlorine, all with a single bond to carbon, and three
lone pairs around the chlorine.
OK…ONE more example, rather more involved, then we'll move on. What is the
structure of CH3COOH? 4 ```````` There are several things to note here. First, in step 2, notice that I bonded carbon with
itself. It warrants repeating that only carbon will do this, and typically, whenever we
have more than one carbon, they will. There are two things to note in step 3. First,
notice that I am starting with only single bonds, and second, I placed both oxygens on the
second carbon. This is because the formula itself gives us the hint that both oxygens
seem to be on the second carbon. However, it would have been equally correct to place
one oxygen on each carbon. The only error would be to place the oxygens on one
another. Also, notice that if I placed both oxygens on the first carbon, or the second
oxygen below the structure rather than on top, all of these and more configurations would
be equivalent, actually the exact same thing, because the Lewis dot structures to not
speak to the actual shape of the molecule. We only would have had a different molecule
if we placed one oxygen on each of the carbons, rather than both to the same carbon.
In step 4, I placed a double bond between the carbon and the oxygen. I could
have just as easily placed it between the two carbons. So, by placing the oxygen on
different carbons, or placing the double bond between the carbons, there are several
structures possible. All of them will fulfill the octet rule beautifully, all of them are
acceptable structures. We call these isomers. Isomers are two (or more) chemicals with
the same formulas (all CH4O2), but different because of different bonding arrangements.
Consider the following Lewis dot structures: 5 ```````` The two structures on the left are identical; each tells us that there are two oxygens
bonded to only one carbon, one with a double bond, the carbons are single bonded to one
another and the three hydrogens are all bonded to the second carbon. They are in
different orientations, so they look different, even though they are the same. The
structure on the right is an isomer; it is actually a different compound, because the
oxygens are bonded one to each of the two carbons, and there is a double bond between
the carbons, rather than just a single bond. This is why the structure on the right is an
entirely different structure.
Sometimes, we end up dealing with ions rather than molecules, such as the polyatomic
ions. In this case, we have to either add or subtract electrons as necessary, one electron
for each magnitude of charge. I personally like taking care of this at the beginning.
This would be easier to understand with a couple of examples. Let's determine
the electronic configuration of Ammonium, NH4+. We notice that with the +1 charge, we
need to remove one electron (recall electrons have a negative charge; to get a +1 charge,
we need one less electron). But from where does the electron come? Well, we can
always try different locations if need be, but usually it doesn't really matter. Let's try
taking it from the Nitrogen. 6 ```````` We'll do a couple more examples of polyatomic ions in the next section.
Extended Octet Rule:
Eventually it became an empirical observation that there are those elements that, for some
reason, could take more than eight electrons in their valence shell. So, they developed
the "extended octet rule", that states that, basically, if we really really really need to, we
can split up lone pair electrons to accommodate more than eight electrons in some
Sound hokey? To me, too, but for a long time they simply could not explain these
deviations. Now, with quantum mechanics, we have come to understand what is
happening. Remember all of those unused orbitals that show up in electronic
configurations for elements past the second row? For instance, we get to the noble gas
Argon, but the 3d orbitals are still empty. Well, those are all ground state electronic
configurations. Turns out, if we go to an excited state electronic configuration, a lot more
We say we "promote an electron" from its "p" or "s" orbital to an empty "d"
orbital. In doing so, we can get more bonding electrons in the valence shell. 7 ```````` Notice that because we are breaking up pairs of electrons, each time we do this, the
number of electrons around the central atom goes up by two. In the above example, this
translated to two atoms, but this may not always be true, because we may also form
multiple bonds. For instance, only one more oxygen could have gone around instead of
two fluorines because oxygen needs two more electrons, and therefore will form a double
bond. Now, we won't have to do this kind of quantitative analysis each time we come up
with one of these problems, but we will have to recognize if an element can undergo
extended octet rule bonding, and if it can, we have to recognize that it will always step up
by two bonding electrons for each lone pair we break apart.
So how do we know when we can, or when we should, use the extended octet
rule? Well, this is how;
1. NEVER use the extended octet rule unless it is absolutely necessary.
2. We can use the extended octet rule only when there are available empty
orbitals. This means any element beyond the second period could qualify.
OK, so, how about one last example? So, what is the Lewis-Dot structure for
sulfate, SO3-2? Well, first of all, since we don't want to bond oxygens to themselves, the
sulfur must be the central element. Furthermore, sulfur is a third row element, so it does
have the necessary empty d orbitals necessary for the extended octet rule. Since it has a 8 ```````` 2- charge, this implies that the polyatomic ion has picked up two additional ions, which I
will add to the sulfur initially. I hope this was helpful to get you started, but in the end, Lewis dot structures are
puzzles, and like any other puzzle, the only way to really get good at it is practice
practice practice. It is an important skill, though, so be sure you can get it down pat. Call
me, email me or stop by to see me if you have problems or questions! 9 ...
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- Spring '10