hello everybody.
today's presentation covers electron configurations
and all things related to electrons.
this is extremely important to the field of
chemistry because the number and position
of electrons in the atoms of different elements
are going to determine their chemical properties.
the periodic table is divided into four blocks
and these are very important for understanding
electron configuration.
the four blocks are the S block, P block,
D block, and F block.
as we look at the next slide, we'll have you
label your individual periodic table also.
this image shows the different blocks of the
periodic table.
the four blocks again are the S block, the
D block, the P block, and the F block.
please note that helium, located here, is
considered to be part of the s block.
to help you remember the different blocks
of the periodic table, try this mnemonic device.
this portion of the periodic table looks like
a mountain, so we'll call this the summit
for the s block.
this other region, the P block, looks like
another mountain so we'll refer to this as
the peak. this area of the periodic table
looks like a valley; D for Death Valley. and
then finally the F block down here looks like
a train track, so we can think of the freight
train that would roll along the train tracks.
take a moment to label your individual copy
of the periodic table, just to the left of
hydrogen you should label 1s. just to the
left of lithium, 2s. just to the left of sodium,
3s, and so on for the s block.
just to the left of scandium, please label
3d. just to the left of yttrium, Y, label
4d. just to the left of boron, label 2p, and
just to the left of aluminum, label 3p. just
to the left of lanthanum, which is right here,
label 4f, and just to the left of actinium,
label 5f. so what do S, P, and D, and F mean?
well these are referring to different sublevels
that exist within a principal quantum level.
this is referring to an energy level.
so if we are looking at the first energy level,
n is equal to one, there is only one sublevel;
it is an S sublevel. for the second energy
level in atoms, there are 2 sublevels.
now we would have an S and a P sublevel being
possible.
in the third energy level, there are now 3
sublevels; S, P, and D. and in the fourth
energy level, there would be S, P, D, and
F sublevels present.
this slide shows a Bohr model of a Neon atom.
please note that the nucleus would contain
10 protons and 10 neutrons. in the first energy
level, principal energy level 1, there would
be 2 electrons and these would be found in
the S sublevel. in the second energy level,
there would be held a total of 8 electrons;
2 of them would be in the S sublevel, the
remaining 6 would be in the P sublevel of
the second energy level.
this slide shows what the different sublevels,
S and P sublevels, would look like.
the 1s orbital is also the 1s sublevel.
this is a spherical region of space, which
surrounds the nucleus of an atom where we
have a high probability of finding 2 electrons.
the 2s orbital or 2s sublevel is again a spherical
region of space, now further out from the
nucleus, which would be a tiny dot at the
center of this image.
the 2p sublevel, actually contains 3 different
orbitals.
the individual P orbitals look like 2 balloons
tied together.
we see the 1 P orbital here, we see another
one illustrated here, and we see another one
illustrated here.
these are going to line up along the 3 axes;
x, y, and z. this slide details the number
of electrons which can be held in different
orbitals or different sublevels.
the S sublevel can hold a total of two electrons.
P sublevels can hold a total of six electrons,
because there are three orbitals, each which
can hold two electrons.
the D sublevel with five orbitals could hold
up to ten electrons, and the F sublevel with
7 orbitals could hold as many as fourteen
electrons.
please note the 1st energy level with only
an S sublevel can hold only 2 electrons.
the 2nd energy level contains both S and P
sublevels, so it can hold a total of eight
electrons - 2+6.
the 3rd energy level with S, P, and D sublevels
can hold a maximum of 18 electrons Because
it contains an S sublevel, a P sublevel, and
a D sublevel.
the fourth energy level could hold up to 32
electrons.
this slide shows the appearance of S, P, D,
and F sublevels for GT chemistry you should
be most familiar with S sublevels and P sublevels.
to determine the electron configuration of
arsenic, an element that contains 33 electrons,
please consult your periodic table.
this is going to be as easy as counting from
1 to 33.
start with element number one, thats hydrogen.
you’ve labeled that on your periodic table
1s, so we’re going to count 1s, and then
1s2. now we’re at number three which is
lithium, so thats 2s, so we’re going to
count 2s1, 2s2, now we’ve counted up through
number four.
next to number five, boron, is labeled 2p,
so we need 2p6, and we’ll continue this
up through number 33, so we get 3s2, 3p6,
4s2, 3d10, 4p3. the first number in electron
configuration is telling us the energy level
which different electrons are found in.
these two electrons are in the first energy
level.
these ten electrons are going to be located
in the third energy level.
additionally, the third energy level contains
six electrons and a P subshell, and two electrons
in the S subshell.
the letter is telling us which block or which
sublevel are the electrons located in.
the superscript is going to tell us the total
number of electrons in that particular orbital
or subshell type. for example, in the fourth
energy level the P sublevel contains only
three electrons so it is half filled.
the diagonal diagram is a good way to remind
yourself which different orbital types or
sublevels will fill in which particular order.
first filled is 1s, followed by 2s followed
by 2p then 3s, then 3p, 4s. after that 3d,
followed by 4p, then 5s, then 4d, 5p, 6s,
4f, 5d, 6p, 7s, 5f, 6d,7p. the noble gas configuration
is a shortcut way to write the electron configurations.
GT chemistry students will need to know how
to do both the full version and the shortcut
version of electron configurations.
please note that the noble gases are going
to be found in the rightmost column on your
periodic table.
they are helium, neon, argon, krypton, xenon,
and radon. also note that each noble gas finishes
a row or an energy level.
this means that the highest energy level in
all the noble gases is going to be 100% filled.
noble gas electron configurations are going
to take advantage of this by condensing what
we need to write in order to convey information
about the electrons that an atom has. for
example, the electron configuration of helium
would be 1s2. the full version for carbon
would be 1s2 2s2 2p2. the noble gas shortcut
for carbon would be to write helium rather
than the 1s2 then fill in the information
that comes after that.
now, for this particular element, the shortcut
does not save us a lot of writing.
let’s take a look at an example where it
does save us quite a bit of time.
earlier we determined the electron configuration
of arsenic, which is found right here; 1s2
2s2 2p6 3s2 3p6 4s2 3d10 4p3. notice that
the part thats in yellow is going to be the
exact same as what the electron configuration
for argon would be.
so to write the noble gas configuration for
arsenic, we would fill in for this inforamtion,
[Ar] in square brackets followed by the information
that’s in purple for us; 3d10 4p3. so this
can save us quite a bit of writing as we’re
writing electron configurations.
the Aufbau principal tells us the order in
which different orbitals will be filled by
electrons as we build atoms.
now please note that each electron is going
to occupy the lowest energy level first.
this is because, well essentially, electrons
are lazy.
so if we have an atom that has only one electron,
that electron would go into the 1s orbital.
with 2 electrons, both can go into the 1s
orbital.
an atom that has three electrons, would place
two in the 1s, and one electron in the 2s.
four electrons would place two in the 1s,
two in 2s. if we have to go beyond that, we
would place one electron into the 2p sublevel
into one of the three P orbitals. and as you
see, we are going to fill up through 2p and
into 3s, and then 3p, and then to 4s, and
you should note that the order that these
orbitals are being listed here, does correspond
exactly with the diagonal diagram.
what is this referring to, is saying that
these are all equivalent energy it means that
the three orbitals here, the three boxes that
are part of the 2p sublevel are all of equivalent
energy.
the three orbitals, the three boxes that are
part of the 3p sublevel, are all equal energy.
the same goes for the five different orbitals
that are part of the 3d, or 4d, or any of
the D sublevels. and same goes for the seven
orbitals which make up either of the F sublevels.
the pauli exclusion principle tells us about
the behavior of electrons within a single
orbital.
these electrons are going to have what we
call ‘opposite spins’. in chemistry we
call them spin up and spin down.
i like to think of tops however; one spinning
clockwise and the other spinning counterclockwise.
in this illustration the box is representing
an orbital.
the up arrow is representing one electron
with an up spin.
the down arrow is representing a second electron
with a down spin.
Hunt’s rule tells us that single electrons
with the same spin must occupy each equal
energy orbital before additional electrons
with opposite spins can occupy the same orbitals.
what the heck does this mean?
it mean that this is not an okay way to write
an orbital notation for nitrogen.
this one is.
why is this the case? i like to think that
these three boxes right here as being seats
on the bus, and because people are not necessarily
friendly on a city bus, for example, if theres
an open available seat people are going to
fill into an empty seat before two people
will sit next to each other.
Hunt’s rule tells us that the orbital boxes
for orbital notations should always be written
following this pattern.
with one, just a single up arrow. two, we
go up up.
three would be up up up.
four we would be up up up and then a single
down.
five would be three ups then two downs. and
then six, three ups first then three downs
following.
this is an orbital notation for the element
V, vanadium.
what we can see here is that we have different
energy levels.
energy level 1, level2, level 2, level 3,
level 3, level 4, and also again level 3.
we also have different types of subshells.
S sublevels, P sublevels, D sublevels.
orbitals are single boxes.
application of the Aufbau principle, Pauli
exclusion principle, and Hunds rule allows
us to number the electrons.
this is important for determining quantum
numbers.
so lets number the electrons.
here wo go.
number one is the up, number two is the down.
3 up, four down.
here's five is the up, six is the up, seven
is the up, 8, 9, 10, at the downs 11 up 12
down 13 14 15 16 17 18 19 20 21 22 23. this
slide shows examples of lewis structures for
the elements of period 2.
please note that a lewis structure tracks
only the number of electrons in the highest
energy level.
we call these valence electrons.
lithium has one valence electron, beryllium
2, boron 3, carbon has 4, nitrogen 5, oxygen
6, fluorine 7, and neon 8.
now as we look down the columns for these
different elements, everything thats in the
same column as neon will have 8 valence electrons
with the exception of helium. all of the halogens
underneath fluorine will have 7 valence electrons.
all of the members of the oxygen family or
group will have 6 valence electrons, and the
same goes for the other columns as well.
please review this information about quantum
numbers.
we will have a class period dedicated to students
writing orbital notations and using that information
to determine the quantum numbers for any given
electron.
this final slide relates electron configurations
and orbital diagrams with information previously
learned about ions. please remember that ions
are charged particles and ions can have either
positive charges or negative charges.
positively charged ions with lose a number
of electrons which is equivalent to the number
of valence electrons they possess.
lithium would have one valence electron as
an atom, it would have an electron configuration
of 1s2 2s1. when it loses that one electron
in the second energy level, it drops having
a full first level and this stabilizes the
particle and causes it to have a positive
1 charge.
negatively charged ions will gain electrons
to stabilize those particles.
the chlorine atom gains a single electron
moving from 3p5 to 3p6 to become stable.
sulfur will add two electrons to completely
fill its 2p - sorry, 3p level to become completely
filled and stabilize the particle.
