In this video we will look at Lewis structures. We will determine the Lewis structure for
an ion, the Lewis structure for some simple molecules, and determine the Lewis structure
for a molecule where the element in the center of the molecule does not follow the octet
rule. Let's start by looking at chloride Cl-. We know that chlorine is in the seventh column
and therefore has seven valence electrons. So we start by saying that we have seven valence
electrons. Now the chloride ion has a negative charge. That's one more electron. We take
our seven valence electrons and our one for the negative charge we have eight total electrons
around the chloride ion. Put the symbol in the middle Cl. Start by putting the electrons
around the Cl one at a time. Now that only gives us four electrons. So now we start to
pair up. You'll notice that we have four pairs around our chloride ion. This gives us an
isoelectric effect of a noble gas. This follows the octet rule. Now let's look at the Lewis
structure for a simple molecule, methane, CH4. In this one we're going to have five
atoms carbon and four hydrogen. Let's start by drawing the symbol for carbon and let's
put the number of valence electrons around carbon. Carbon is in the fourth column of
the periodic tables so it has four valence electrons. We start by putting one electron
around carbon then two three and four until we have all four valence electrons. Notice
that we did not pair any of the electrons. We then look at hydrogen. Hydrogen is the
first column so it only has one valence electron. So we'll put one dot, one electron, next to
the hydrogen. We do this for all four hydrogens. To come up with a Lewis structure for this
molecule we're gonna have to combine these and come up with bonding pairs. For there
to be a chemical bond we have to share two electrons so we'll start with a carbon in
the center. It has the most nonbonded electrons then we'll start attaching our hydrogens to
these noticed that now one bond from carbon, one bond from hydrogen we now have a bonded
pair. Notice that carbon has four hydrogens around it. This is the Lewis structure for
methane. Also notice that there are no lone pairs on the hydrogen because each hydrogen
only had one electron. Our next example let's look at carbon dioxide. In carbon dioxide
CO2 we have carbon with its four valence electrons just as in methane, oxygen in the sixth column
of the periodic table will have six valence electrons. Again draw the first four electrons
unpaired. Now electron five and electron six we're gonna have to pair up. This gives us
two pair sets on our oxygen. These are lone pairs and will not be bonding unless necessary
with the other atom. Now we put carbon in the center again because the carbon has the
most unpaired electrons. Put an oxygen on either side of the carbon and notice we now
have two bonded pairs. We should check our octet when we count we see that carbon only
has six electrons and both oxygens only have seven. This is not sufficient for the octet
rule. So we're gonna move more electrons in. We're gonna move an electron from the carbon
and one electron from the oxygen between the first carbon and oxygen and do the same with
the second oxygen and the carbon. Now notice we have two bonded pairs between the carbon
and the oxygen. This would be a double bond. So the Lewis structure for carbon dioxide
has double bonds between two oxygens. When we double check our octet notice that this
Lewis structure has eight electrons around the carbon and both the oxygens. This is the
correct Lewis structure for carbon dioxide. For our list example let's look at a molecule
where the central atom does not necessarily follow the octet rule. A good example of this
is aluminum or boron. Both of these can exist with stable molecules with only six electrons
instead of the normal eight that follows the octet rule. Another couple of examples where
you'll have stable ions or compounds where you do not follow the octet rule is phosphorus
and sulfur. Both of these often are hypervalent meaning that they have more than eight. Such
is the case is in sulfate. Let's look at a Lewis structure for our aluminum chloride
AlCl3. Here aluminum being in the third column only has three valence electrons so draw the
three electrons around the aluminum represented by dots. Remember not to pair and chlorine
we've already shown chlorine with seven valence electrons. We show those three chlorines and
then we start to combine. Notice that in this case we have a chlorine on bonding with each
one of the electrons from the aluminum and on the aluminum we have an empty spot. We
will find out as we move on through chemistry that this is actually an empty P orbital but
at this point because we know that aluminum is stable without following the octet rule
we can leave this empty. Lewis structures are very important. We're gonna use them over
and over again. It's important that you practice with Lewis structures and that you can take
a simple chemical formula, using Lewis structures, the number of valence electrons, and charges
if there are any and come up with a stable structure. Practice and get confident.