[MUSIC]
Lighthouse Scientific
Education Presents,
a lecture in the
Atomic Structure series.
The topic; 'Valence Electrons,
Octet Rule and Dot Structure'.
Material in this lecture builds
upon the previous lectures
Quantum Mechanics
and Electronic Configurations
This lecture looks at three
useful tools that follow
from the electronic
configurations of the atom;
Valence Electrons, we will
determine them directly
from electronic configurations
and from the Periodic Table.
We will use valence electrons
as the basis of the Octet Rule
and in developing electron,
or Lewis, dot diagrams.
We start by revisiting the
stratified energy level diagram
to state that electrons
in the UPPERMOST occupied
energy levels of the atom
are the ones that participate
in chemical reactions.
In this lecture we are taking
the lessons we learned
in determining electronic
configurations to prepare
ourselves to transition
from discussing the
atom to considering how
atoms  form compounds
and molecules.
Real CHEMISTRY!
For the atom represented
by this diagram, the
electrons in the 3s and
3p orbitals are the ones
in the uppermost
occupied energy levels.
But why just concern
ourselves with these orbitals?
Well, the uppermost
most orbitals have the
highest energy electrons.
 Also these electrons
are the furthest away
from the nucleus and
are therefore the
least tightly held.
The combination of being
of a higher energy and held
with a loser grip make
the upper orbital electrons
natural candidates for
the sharing, borrowing and
releasing of electrons that
define chemical reactions.
And those elections located
under the upper orbitals?
If we look at the electronic
configuration of just those
electrons, we see
that they have the same
electronic configuration
as their nearest noble gas
(neon in this case).
And, as we know, that
configuration is
particularly stable.
Let's take a closer look at
this. Once an energy level is
completely filled
those electrons
"do not" participate
in chemical reactions.
There are quotes around do
not because things in reality
are always more complicated
then they first seem.
However, we will stay with
the 'do not' in this lecture.
So, where are the  energy
levels, as designated by the
principle quantum
number n, complete?
That's right it is
at a noble gases.
n =1 is complete at helium; n=l
only has an s orbital and 1s2
is that orbital filled.
A filled energy level can
use the bracket short hand,
brackets helium
is equal to  1s2.
For n = 2,   the energy
level has 2s and 2p orbitals.
The noble gas neon has the
right number of electrons
to fill those orbitals.
Brackets Ne
If we replace 1s2 with brackets
helium we can more clearly
see that the outer occupied
orbitals are 2s and 2p.
The same argument can
be made with argon.
The 3s and 3p orbitals are
the outer occupied orbitals.
They come after the filled n =2
energy level of neon. When the
energy level is 3 or larger
there is a complicating matter.
These higher energy levels
have additional orbital types
like d and f that require
additional consideration.
But for now we will just
focus on electrons in the
s and p orbitals. Electrons
in the s and p orbitals,
after the last noble
gas configuration,
are called the valence
electrons; ve for short
Valence electrons are a
very important concept in
considering chemical reactions
and molecular structure.
It is worth our time to make
sure that we are confident
of which elections in an
atom are characterized as
valence electrons. Presented
here is a list of the
first 18 elements along with
their electronic configurations.
We are going to note
the valence electrons
for each element.
Revisiting the definition:
electrons in the s and p
orbitals, after the last noble
gas configuration, are
called the valence electrons.
To be clear we will highlight
the noble gas configurations.
Helium is the first noble
gas and both of its electrons
are called valence
electrons even though
they are technical not after a
noble gas. Then there is neon.
It has a total of eight
2s and 2p electrons
which come after the noble gas
configuration of helium.   And
finally, there is argon. It has
a total of 8 3s and 3p electrons
which come after the noble
gas configuration neon.
The pattern two-eight-eight
is an important arrangement
for these first 18 elements.
It is a recognition that it
takes the first 2 elements
before the n=1 energy level
is filled. And 8 more elements
to fill the s and p orbitals of
the n=2 level. And an additional
8 more elements before
the s and p orbitals
are filled in the
n=3 energy level.
We will see this pattern,
2:8:8, again and again.
For instance if we look at
the Periodic Table, it shows
us that there are 2
elements in the first row.
Helium has been moved
to its other position,
the one next to hydrogen
for this demonstration.
There are 8 elements
in the second row
and 8 in the 3rd row.
Okay, now to listing out
the valence electrons
for each element.
For elements of the 1s block;
Hydrogen's lone electron
will be considered
a valence electron.
Helium as mentioned
has 2 and that caps the
valence electrons
in the n=1 level.
We will no longer be
considering electrons from
the 1s orbital.
It is full
For the 2s block;
Lithium has 1 electron
after the noble gas
configuration of helium.
Beryllium has 2 electrons
after the noble gas
configuration of helium.
Now, including the
2p orbital block;
boron has 3 electrons after
the noble gas configuration
of helium. Two are from the
2s block and 1 from the 2p.
Carbon has 4; two from
the 2s and two from the 2p.
Nitrogen has 5: two from the
2s block and three from the 2p.
Oxygen has 6; two from the
2s block and four from the 2p.
Fluorine has 7; two from the
2s block and five from the 2p.
And, as stated neon has 8; two
from the 2s block and six from
the 2p. We will no longer be
considering electrons from the
2s and 2p orbitals
because they are full.
For the 3s orbital block.
Sodium has 1 valence
electron after the
nobel gas configuration
of neon. 3s1.
Magnesium has 2
valence electron
Now including the 3p block.
Aluminum has 3
valance electrons.
Two from the 3s block
and one from the 3p.
Silicon has 4.
Phosphorus has 5.
Sulfur has 6.
Chlorine has 7 and argon
with the stable 8 arrangement;
two from the 3s block
and six from the 3p.
Viewing just the part of
the electronic configurations
with valence
electrons we see that
as we go up by one
in atomic number
we go up by one in
valence electrons
until we reach a noble gas.
Then there is a reset to 1
and the incremental
increase begins again.
There is another, and
probably easier way,
to obtain the correct number of
valence electrons and that is to
use the Periodic
Table of Elements.
Look at your Periodic Table.
Notice that at the top of each
column in the s and p blocks
there is a roman numeral
followed by the letter A.
The A stands for
main group elements.
The roman numeral tell us
how many valence electrons
the elements in
those columns possess.
Now some Tables don't use
the roman numeral system.
They just number the
columns from 1-18.
We will relate the two
styles as we go along.
Look at the first
column on the left.
It has the roman
numeral I or 1.
These elements have 1 ve.
On top of the second
column there is a
2. These elements have 2 ve.
Hopping over to the p block
then there is 3, 4,
5 (as a V), 6 7 and 8.
Each is a count of
the ve in that column.
Again, helium has been
moved to its other position,
the one next to hydrogen
for this demonstration
since it has 2
valence electrons.
The roman numeral style and
the regular number style at
the top of the columns are the
same for the 1st two columns.
But when we go to the
start of the p block
the roman numerals
continues on at 3
but the regular numbers
jump 10 values to 13.
Reasonable considering
that they were counting the
10 columns of the
transition metals.
It is easy to reconcile
the two systems by
subtracting 10 from the
regular number style. That is
13 represents roman numeral
3 and its 3 valence electron,
14 represents 4, 15 5 and
so on to 18 representing
roman numeral 8 and
its 8 valence electron
To restate the use
of the Periodic Table
in finding valance electrons;
the elements in the first
column have 1 valence electron.
Moving across the Table, the
elements of this column have 2
valence electrons;
these have 3, 4, 5, 6, 7
and the noble gases
have 8 valence electrons.
Using the Periodic Table is
an easier way to get the number
of valance electrons but
being able to get them using
an electronic
configuration shows a deeper
understanding of what
valence electrons really are.
It's best to be aware
of both procedures.
A little quiz
should see to that.
Given carbon's
electronic configuration,
how many valence
electrons does it have?
We see that after the filled 1s
orbital it has two electrons in
its 2s block and two in its 2p.
That's 4 valence electrons.
Consulting the Periodic Table
and finding carbon, we note
that the roman numeral
on the top of the column
also gives us 4 as does
subtracting 10 from 14.
What about fluorine?
Consulting its electronic
configuration we see that
it has two electrons in its 2s
orbital and five in its 2p block.
for a total of 7
valence electrons.
Using Periodic Table
yields the same value of 7.
One more, Silicon
Looking at the electrons that
come after the last noble gas
configuration we find it has two
electrons in the 3s orbital and
two in the 3p block of orbitals.
Summing to 4 valence
electrons. Not surpassingly the
value from the Periodic Table
also provides us with 4.
Even though carbon and silicon
are in different rows
they have the same
number of valence electrons.
The reason we spent so much
effort on determining the
number of valence electrons is
that we are going to need that
number. The first place we will
apply it is with the octet rule.
And that rule says that
atoms with 8 electrons
in their outer shell
(and we are only
talking s and p orbitals)
are the most stable.
The 8 comes from filling the s
orbital, two electrons, and the
three p orbitals (each holding
2 electrons for a total of 6);
2 plus 6. Noble gasses
have an octet (s2 p6).
One can also see that they
fall under the 8A heading.
The Octet rule mostly applies
with elements in the 2nd
and 3rd row. A duet, or 2,
rule can be used for row 1:
Helium is a noble
gas with 2 electrons.
And as a heads up
for chemistry to come,
'atoms interact with
other atoms such that each
has 8 valence electrons (like
their nearest noble gas)'.
There are some limitation to
this statement but it holds for
a lot of chemistry
covered at this level.
It's worth sketching out a
model of the atom consistent
with the octet and duet rules.
Using a simplified description
of the atom and  orbital shells,
we can generate a
visual model of the atom
that can help us better
understand chemistry topics
like atom-atom interaction,
ion formation and,
conveniently, will
lead to our next topic
Lewis diagrams.
Let's see how that model works.
We begin by drawing the
positively charged nucleus as
a small sphere. Starting with
the duet rule and the n=1 shell.
We can represent the first
shell as a circle or sphere
around the nucleus.
Duet means 2.
Our first shell is complete with
2 elections; paired  electrons
Two dots are added to the
first shell as a representation
of the paired electrons.
Next we move to rows 2 and 3
which follow the octet rule.
Row 2, with its 8
elements, gets a shell
that can hold up
to 8 electrons.
Burrowing Hund's rule, which
has opposite spin elections
pairing, we add pairs of
dots to the second shell
and we add them as far
away from each as possible.
Not a bad idea since
negatively charged electrons
repel each other. Therefore
there are 4 sets of dots.
For row three and its
8 elements we add a
new larger shell which
can also hold 8 electrons.
As with the second shell,
these electrons are paired
and the pairs are placed
as far apart from
each other as possible.
This is our octet (rule)
version of the atom,
or at least atoms having
18 or less electrons.
Still, it is only a model and
is best used as a bookkeeper
that keeps track of
valence electrons.
Speaking of 18 electrons,
the pattern of electrons in
this atom with full shells is
that same 2:8:8.    That pattern
comes up regularly in
the study of chemistry
because it represents
the maximum number of
valence electrons in
each shell. 2 in the first
and 8 in the 2nd and 3rd.
It also helps us explain
the formation of ions and it
leads us to the next topic:
The electron dot diagram
or, Lewis Structure.
This a useful way to
visualize the valance electrons
around an element (or an atom).
It is very similar
to the model of the
atom we just constructed
with the octet rule
but it is only concerned
with the last or the
outer shell electrons.
Let's begin with an overview.
The symbol for the element
is written in the center.
One dot is added in
'a specific pattern'
for each valance electron.
Along with the octet rule
and duet rules the
electron dot diagram
can be used to help
understand how atoms form
molecules and ions. But that's
a lecture for another time. First
we need to be able to construct
diagrams for just atoms.
Determining the Lewis dot
diagram of an atom is similar
to that of filling its
electronic configuration.
It is not as accurate as
an electronic configuration.
Like the atom model
using the octet rule,
dot diagrams are practical
tools. We are going to
construct a platform for
generating dot diagrams.
We will use a capital
E as a generic stand in
for all symbols of elements.
We will use dots
to represent valence
electrons and we give element E
six valence electrons for
demonstration purposes.
The 'dot' diagram or 'dot'
structure can be considered as
using 4 equal energy orbitals.
Loosely, we can view this
as putting the single s
orbital of an energy level
on the same footing
as the three p orbitals
giving us 4 equal
energy orbitals.
And, as we did in the with
the model of the octet rule,
place these 4 orbitals
(pictured as boxes) around
the symbol of the
element and put them
as far away from each
other as possible.
Our next consideration
is how to add electrons to
the boxes. For that we
will consult the Hund rule
and the Aufbau principle.
The Hund Rule limits these
"orbitals" or boxes to
no more than to 2 electrons.
We can generate the
Hund rule on the diagram
by dividing each box in half.
There are 8 half boxes,
The format of the dot
diagram is complete.
Now to the pattern of
adding electrons as dots.
According to the Aufbau
principle, electrons are added
to equal energy
orbitals before pairing.
Remember how, when deriving
electronic configurations,
we added one electron to the
2px then another to the 2py
then another to the 2pz orbitals
before any pairing occurred?
Same thing applies here.
One electron is added
to each of the equal energy
orbitals before any pairing.
That uses up 4 of element
E's 6 valance electrons.
It is only after
each box has a dot
that we continue
assigning dots by pairing.
The additional 2 valence
electrons are paired.
A couple of small
details to clean up.
We do not actually have
to include the boxes around
the element. They were used
here to develop the concept.
Also, since the boxes
represents equal energy
orbitals. it is not strictly
incorrect to start with the dot
in a box other
than the one on top.
It is, however, by convention
that we start at the top box
and go in a clock wise
matter around the symbol.
And that's the basics of the
dot diagram.   We will use some
basic rules for constructing
Lewis Dot Diagram
until we feel comfortable
with the procedure.
Begin by writing the symbol
for the element and yes,
we will place it at the center.
Determine the number of
valance electrons for the
element. This can be found
by using the electronic
configuration or
the roman numeral
or column number
above the column on
the Periodic Table
where the element resides.
Now, add electrons, as dots,
to each box (that's an orbital)
around the element
symbol (up to 4 ve).
And do this in a clockwise
fashion.  If there are more
than 4 electrons, repeat
the pattern of addition
to each box and pair.
Shall we give it a try?
We start by writing the
symbol of the element.
Carbon has offered to go first.
Find the number of valence
electrons for carbon.
Referring to its electronic
configurations we see that
electrons coming after the
last noble gas are 2s2 and 2p2.
2+2=4. Carbon has
4 valence electrons.
The dot diagram
will have 4 dots.
Next step is to add dots,
up to 4, around the element,
in a clockwise
manner. 1, 2 3, 4.
Note we have added one
dot to each of our boxes
without pairing and that uses
all of our valence electrons.
This is carbons dot diagram,
So, the last step about pairing
electrons isn't required here.
We should do another
one, step by step.
How about fluorine?
We start by writing
the element symbol, F.
We need to find the number of
valence electrons for fluorine.
According to its electronic
configuration there are
2 valence elections in the
2s and 5 in the 2p. That's 7.
We could have also located
fluorine on the Periodic Table
and note the roman numeral V I
I, or 7, on top of the column.
Column number 17 minus
10 would also yield 7.
Fluorine has 7
valence electrons.
The dot diagram
will have 7 dots.
Now we add the electrons
as dots around the F
one at a time, starting at
the top, until we get to 4.
1, 2, 3, 4. That's 4 of 7.
We still have 3 more dots.
Step 4 has us repeating
the pattern, adding dots one
at a time. These dots, however,
will pair.  Starting at the top
5, 6, 7. Three boxes are
filled and one is half-filled.
All right, now we are
going to run through the
first 18 elements and
generate dot diagrams for each.
We are attempting to develop
the pattern of using the rules
so that we do not have to
relying on using the rules.
To speed matters up we
will use a table that states
the number of valence
electrons for each element.
It's ok to use a table.
Starting with hydrogen and
its one valence electron.
It's one dot goes
in the top box.
Helium, 2 valence electrons.
One on top, two on top.
Hmmm, looks like the noble
gas helium is going to follow
the duet rule.
It will not have
4 boxes around
the element symbol.
It will have one box, or orbital,
on top and the dots will pair.
Ok, we'll just going to
have to remember that.
On to lithium with
1 valence electron.
Remember, valence
electrons come after the
last noble gas configuration
and that is helium's 1s2.
We will put that dot on top.
And we see that lithium
and hydrogen have
the same dot diagram.
Makes sense since they are
both under the IA heading.
To beryllium, it has
2 valence electrons.
One on the top, the second
added in a clockwise fashion.
Boron, 3 valence
electrons.  1, 2, 3
Carbon, it has 4
valence electrons.
1, 2, 3, 4. All are added
in a clockwise fashion.
And each of the
boxes has 1 dot.
Anymore and we
will have to pair.
Which is the case with the 5
valence electrons of nitrogen.
The first 4 go in
clockwise fashion
and the 5th pairs
in the top box.
Dots added after the 1st 4
dots repeat the pattern of
going around in a
clockwise fashion
and that includes
starting at the top.
Oxygen has 6 valence electrons.
1, 2, 3, 4 and pairing 5 and 6.
Fluorine's 7 valence
electron dot diagram will look
like oxygen's but
with one more dot.
1, 2, 3, 4 pairing 5, 6 and 7
Neon, the noble gas
has a full octet.
That means 8 valence
electrons and 8 dots.
4 dots around and
4 more in pairing.
There is no room for more
dots (that is electrons)
on this element and that is why
full octet atoms are so stable.
It also means that as we
move to the next element
we will be starting
on a new shell
and starting back at
1 valance electron.
Row 3 on the Periodic
Table starts with sodium.
We are at the beginning of a
new set of valence electrons
(coming after neon).
Sodium has 1.
And it goes on top like
hydrogen and lithium.
Each element is a
member of column IA
Magnesium has 2 valence
electrons  1, 2 but so as does
beryllium. They both
reside in column IIA.
Aluminum with its 3
valence electrons shares the
same dot structure as boron
as well as sharing a
column on the Table; IIIA
Silicon has 4 valence
electron. as does carbon
1, 2, 3, 4. They are
found in column IV
Phosphorous and nitrogen
both have 5 electrons.
1, 2, 3, 4 and pair the 5th
electron. They share the same
dot diagram and have
similar physical properties.
You can find them in column V.
Sulfur and oxygen both
have 6 valence electrons.
Sulfur's dot diagram is
going to look like oxygen.
1, 2, 3, 4 and pair the 5th
and 6th valence electron.
They share the same
dot diagram and have
the same physical properties.
They belong to column VI.
Is the value of the dot
structure becoming clearer?
The Halogens chlorine
and fluorine are both
one valence electron
short of a full octet.
It shouldn't be too
surprising to find that these
atom types readily 'steal' an
electron to fill their octet.
The 7 dots around chlorine
have the same pattern
as the 7 dots around fluorine.
They do, after all, reside
in the column VII elements
And finally, the
noble gas argon.
It does have a full octet
and is a stable atom type
just like that
other noble gas Neon.
Try your own pass through
the 1st eighteen elements
and this important
topic will quickly become
more comfortable and
add as a skill set.
Let's recap what
we have learned
Electrons in the UPPERMOST
energy levels of the atom
are the ones that participate
in chemical reactions.
The ones beneath
them are in stable
noble gas electronic
configurations.
Electrons in the
s and p orbitals,
after the last noble
gas configuration,
are called the
valence electrons (ve).
We use valence electrons
in the Octet Rule:
Atoms with 8 electrons
in their outer shell
(specifically s and p
orbitals) are most stable.
That 8 comes from filling the
one s orbital (2 elections)
and three p orbitals
(6 electrons).
A duet, or 2, rule
can be used for row 1;
That means the 1st 3 rows go
2:8:8 for valence electrons.
The electron (or Lewis) dot
diagram (or dot structure)
is a useful way to visualize
the valance electrons
around an element.
The rules for creating
a dot diagram begin
with writing the symbol for the
element (it will be in the center)
Determine the number
of valance electrons;
either through the element's
electronic configurations
or the column in the Periodic
Table (we can either use the
roman numerals or column
number above the element).
Add electrons, or dots, up
to 4, in a clockwise fashions,
one at a time, to each
side of the element symbol.
We called theses equal energy
orbitals boxes in this lecture.
If there are more than 4
valence electrons repeat
pattern of addition to each
side and pair electrons or dots.
And that concludes our
lecture on valence electrons
and their use in the
octet rule and dot diagram.
This is a useful topic that make
sense with a little practice.
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