Hello I'm The Chemistry Guy. In this video we will learn more about atomic
structure, and in particular learn about the atomic orbitals and how those
orbitals correspond to the Periodic Table.
In the presentation called "Atomic Structure" we learned that electrons are particles
that also behave as waves. This is known as the particle-wave duality, and in fact
all objects that are moving have this particle-wave duality, but the electron
is the only particle that is small enough that this is actually a
measurable amount. The wavefunctions describing the probability density of an
electron can be solved for the hydrogen atom, and these wave functions
depend upon three quantum numbers n, l and m. As we learned earlier, the quantum
numbers also have the following rules: n is an integer one two three etc for
each value of n, l can be 0 up to the number 1 less than n, and for each value
of l, m can be from minus l up to positive l. Again to review, n is called
the principal quantum number and it determines the size of the electron
density contour, which is just called an orbital. As n gets larger the orbital
gets larger and on average the electron is further away from the nucleus.
l determines the shape of the orbital and each value of l has a 1-letter
designation. If l is 0 it's an s-orbital, or spherically shaped. If l is equal to 1
it's a p-orbital, l is equal to 2 it's a d-orbital, and L is equal to 3 it's an f-orbital.
m determines the orientation of this orbital or electron density contour,
and in general there are 2l plus 1 such orbitals for each value of l.
We also learned that electrons have a magnetic moment that kind of like a little magnet,
and we call this the electron spin. It is denoted by the quantum number s.
s can have values of -1/2 and 1/2, which we call spin-up and spin-down.
Therefore the electron in an atom is labeled by four quantum numbers n, l, m, and s.
Finally the Pauli Exclusion Principle states that no two electrons
in an atom, or even in a molecule, can have the same four quantum numbers. this
This means that a specific orbital, which for an atom is a set of allowed values of n, l, and m,
can hold only at most two electrons, one spin up and one spin down.
If an atom only has a single electron the energy of the electron is only
determined by the value of n. This means that the 2s and three 2p orbitals all
have the same energy. As soon as the second electron is added this is no
longer true. Here we will use the Periodic Table to see how the orbital
energies are affected as more electrons are present in the atom. the simplest
The simplest atom is hydrogen, which generally has only a single proton in the nucleus and
a single electron around it. It therefore has an elemental number of 1 and is up
here in the upper left-hand corner of the Periodic Table. This electron will go
into the 1s orbital, so we denote the electronic configuration as 1s1 where
the superscript 1 is the number of electrons in this orbital.
Helium has two protons in the nucleus and therefore has two electrons around it.
The second electron also goes into the 1s orbital; one of the electrons will have a spin up
and the other electron will have a spin down and the electronic configuration is 1s2.
These two electrons are all that will fit in the n equals 1 orbital,
which is called the first shell. The column at the right are all closed shell
atoms and are called the Noble Gases.
Lithium has three protons in the nucleus and three electrons around it in the
neutral atom. The first two electrons again go into the 1s orbital, one spin up
and one spin down, and the third electron goes into the 2s orbital. Therefore the
electronic configuration is 1s2 2s1.
Beryllium has four protons in the nucleus and four electrons; two electrons go into the 1s
orbital and the other two going to the 2s orbital. Again the two electrons in a
given orbital must have opposite spins one up and one down.
The electronic configuration for beryllium is 1s2 2s2. Now remember it is
not a filled shell because you've still got all of the 2p orbitals in that
second shell. That is why it's sitting next to lithium.
Boron has an atomic number of 5. Therefore it has 5 protons in the nucleus and 5 electrons around
the neutral atom. Two electrons go into the 1s orbital, two go into the 2s orbital
and the fifth electron goes into a 2p orbital. Remember the two electrons in
the 1s and the 2s orbitals have to have opposite spin.
The electronic configuration of boron therefore is 1s2 2s2 2p1.
Carbon, atomic number six, has six protons in the nucleus and six electrons around it.
It's electronic configuration is 1s2 2s2 2p2.
Since electrons are negatively charged they repel each other.
Therefore the two electrons in the 2p orbitals will be in different orbitals,
meaning different values of m. There is a slight advantage in
having the same spins; the magnets will be aligned so they will both be spin up,
for example, when aligned with an external magnetic field.
The electronic configuration for nitrogen, which is atomic number seven, is 1s2 2s2 2p3,
where each of the three electrons in the 2p orbitals will be in different orbitals
and will all have the same spin; all spin up or all spin down.
Oxygen, atomic number 8, has 8 electrons around the nucleus so it's electronic configuration is 1s2 2s2 2p4.
Two of the electrons in the p-orbitals will have to go into the same orbital, so one must be
spin up and one must be spin down, and the other two will go into separate p-orbitals
with the same spin. fluorine has nine electrons is atomic number nine
Fluorine, atomic number 9, and nine electrons and it's electronic configuration is 1s2 2s2 2s5.
Two P orbitals will be doubly occupied, one spin up and one spin down,
and the third a 2p orbital will only contain a single electron. neon completes
Neon completes the second row of the Periodic Table. It is atomic number ten so it has ten
protons in the nucleus and ten electrons around it and it will fill all of the 1s
2s and 2p orbitals. it's electronic configuration therefore is 1s2 2s2 2s6.
It is a closed shell system and it
is the second member of the Noble Gases.
We can see that we have a noble gas whenever a set of p-orbitals within a shell, which is a
value of n, is filled. Argon is the third noble Noble Gas.
It has an electronic configuration of 1s2 2s2 2p6 3s2 3p6,
filling the third row. It is considered a closed shell system even
though there are 3d orbitals present that are unoccupied.
So again Noble Gases only really are filling the s and p orbitals of a given shell.
Looking at the 4th row of the Periodic Table, the first two electrons
go into the 4s orbital, the next 10 electrons go into the 3d orbitals, and
the final 6 electrons go into the 4p orbitals. The electronic configuration
for Krypton is 1s2, filling the first shell, 2s2 2p6, filling the second shell,
3s2 3p6 4s2 3d10 4p6.
This is cumbersome so instead we can just write
Ar 4s2 3d10 4p6, where Ar just means the
electronic configuration of argon, which takes care of the first and second
shells and the 3s2 3p6. Things get a little crazy when
we get to the sixth row of the Periodic Table because, as you can see, here we
have a breakout which is the Lathanide series, which is actually represents the
4f orbitals. So as we start to fill across, the first two electrons will go
into the 6ss orbitals, and then we start put one into a 5d and jump down
here, so by the time we get to Lutentium, Lu, which is atomic number
71, the electronic configuration is 6s2 4f14 5d1.
From there we fill up the rest of the 5d orbitals so that by Mercury (Hg) which
is here, atomic number eighty, the electronic
configuration is xenon 6s2 4f14 5d10.
Then we add the next six electrons to
the 6p orbitals to get radon, atomic number 86, and therefore its electronic
configuration is Xenon 6s2 4f14 5d10,6p6.
As you've probably figured out, the Periodic Table
has the following form. Each row corresponds to the quantum number n so
as you go across for the s and p-orbitals they get larger for the outermost electrons for each Element.
The first two columns are for
the s-orbitals. These columns over here are generally for the p-orbitals. In the
middle we've got the d-orbitals but the d-orbitals are for the quantum number
one below the row. So even though this is the fourth row these are the 3d
orbitals here, and then we've got the Lanthanides and Actinides down here that
correspond to the f-orbitals. They are for the shell two lower, so again when
I've got the 6s, I then go to the 4fs,
the 5d's and then the 6p's. The only exception to this periodic table is
Helium. Some Periodic Tables place it above Beryllium since it's electronic
configuration is 1s2 or two electrons in an s-orbital. But since it
fills the n equals 1 shell, most Periodic Tables place it in the last column
so that all of the Nobel Gases are together.
Nobel Gases are very stable and other atoms, if possible, would like to look
like a Noble Gas. Therefore Lithium, which is basically the two electrons in
the 1s plus a single electron in the 2s, easily loses an electron to become
Lithium plus. Remember that's called an a cation, or just an ion, so that it's
electronic configuration is the same as helium. Calcium will lose two electrons
forming Ca 2+ so that it's electronic configuration is the same as Neon.
Losing an electron is easier as the atom gets bigger in other words as the n
or the shell of the last electron gets larger.
At the other extreme fluorine and chlorine would like to gain an electron to become F- and Cl-, or anions,
so that they look like Neon and Argon, respectively. This will fill either
the second shell or the fill out the 3s and 3p electron orbitals. The desire to
gain or lose an electron is called the electronegativity. Atoms with a lower net
electronegativity will lose electrons and those with a high electronegativity
will want to gain electrons. Therefore the electronegativity increases from the
lower left of the periodic table to the upper right,
not counting the noble gases as they already have the desired electronic configuration.
So again the elements down here in this corner have the lowest
electronegativity and will easily lose an electron or two electrons depending
upon what column they're in and it goes all the way up to here with Fluorine and
Chlorine having very high electronegativities because they want to
gain an electron so that they look like two Nobel Gases.
Carbon is in a weird position. it's electronic configuration is 1 s
square 2 s squared 2 P squared therefore It would have to lose 4 electrons from
the n equals 2 shell, or gain 4 electrons, in order to look like a Noble Gas.
Neither of these are possible, so instead it shares its 4 outermost electrons,
called valence electrons, with an electron from up to four other atoms so
that it effectively has eight electrons in this outermost shell. This is known as
the octet rule. We will talk about chemical bonding in "Atomic Orbitals,
Molecular Orbitals and Bonding. As a final point on the Periodic Table, the
Elements are put together into various groups. In the first column with the
exception of hydrogen we've got the Alkali Metals then and they will all
want to lose an electron. In the second column there are the
Alkaline Earth elements, and they want to lose two electrons.
Through here we've got the Transition Metals. We then have the metals. We know that
aluminum is a metal. These are called the Metalloids. Silicon is a semiconductor
used in computer chips because it kind of allows electrons to
flow. Then Hydrogen and this group right here, which is the great majority of the
chemistry in our bodies, are nonmetals they only want to form bonds.
Generally though, Hydrogen can lose an electron. The next group are the Halogens.
They want to gain an electron so that they look electronically like the Noble
Gases. Then down here we've got the Lathanides and we've got the Actinides.
Please note that this picture is not mine, I got it from the web, and they
messed up on the colors between the Lathanides and Actinides. Lathanides should be light
purple and the Actinides should be the salmon color.
I hope you enjoyed this video and if you enjoyed it I hope you will support my
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