Download Arrangements of electrons in the orbitals of an atom is called its

Arrangements of electrons in the orbitals of an atom is called its electron
configuration. And we use the Pauli exclusion principle and other rules to
arrange the electrons. To designate the electron configuration we use the
level number and the letter of the sublevel and a superscript number to
represent the number of electrons contained in the sublevel. Writing the
electron configuration requires that we recall how many orbitals are
contained in each type of sublevel
For example hydrogen with one electron has an electron configuration of
1s1. The orbital diagram for hydrogen can be represented in the following
way.
This notation uses a box to represent the orbital, the label for the orbital
and an arrow to represent the electron. The electronic configuration for
hydrogen can be written as 1s1. This is a short-hand notation which
identifies the level, the sublevel and the number of electrons in the
sublevel. We can also display the energy level diagram for the hydrogen
atom. A portion of the energy level diagram is shown,
So we have three ways to represent the electron arrangement in an atom.
The orbital diagram, the electron configuration and the energy diagram.
All three ways are useful.
The next atom is helium with 2 electrons. So the second electron could go
into the 1s orbital with the opposite spin of the first electron or it could go
into the next orbital in the n = 2 level. It turns out that the energy required
to accommodate two electrons in the 1s orbital is significantly less than the
energy required to place the second electron into the higher energy n = 2
level. The orbital diagram for helium is,
So while hydrogen has the electron configuration of 1s1, helium has the
electron configuration of 1s2. The energy diagram for helium is shown as
here. Notice that there has been a change in the relative energies of the 2s
and 2p orbitals. This is an important point that must be addressed at this
point.
In the hydrogen atom the sublevels in each principle level are degenerate.
In multi-elecron atoms the degeneracy of the energy of the sublevels is lost.
When the third electron is to be placed it must go into the second level. The
first level is filled and can not accommodate any more electrons. When the
electron is added to the second level it can go into the 2s orbital or the 2p,
the question is which orbital is the electron placed? The primary criteria,
the Aufbau principle, states the electrons are to be placed into the orbital
of lowest energy. So we must consider which orbital, when the electron is
placed into it, has the lowest energy? This is answered by considering some
complicated mathematical calculations. The essence of these calculations is
that when an electron is placed into the 2s orbital the electron is likely to
spend more time closer to the nucleus than an electron in a 2p orbital. If
the electron spends more time closer to the nucleus the electron will
experience a greater attraction to the nucleus and it is lower in energy. It
can be stated the 2s orbital penetrates closer to the nucleus than does a 2p
orbital.
So the orbital diagram for lithium is shown below. The electron
configuration for lithium is 1s22s1.
The energy level diagram, on the left shows the relative energy of the 2s
and 2p orbitals based on the ability of the sublevels to penetrate to the
nucleus.
The next element is beryllium which has four electrons. The orbital
diagram for beryllium is shown here. The electron configuration is 1s22s2.
The fourth electron is placed in the 2s orbital. The energy required to pair
the first 2s electron is less than the energy required to place the electron
into the 2p orbital.
The next element is boron with 5 electrons. The orbital diagram for boron
as shown has the one electron in the 2p orbital. The electron can be placed
in any of the three 2p orbitals. The electron configuration for boron is
1s22s22p1.
The energy level diagram for boron is show below.
For the next element, carbon, the sixth electron must be placed in the
correct orbital. The question becomes whether the next electron should be
pair the other electron or whether the electron should be placed in an
empty 2p orbital.
According to Hund's rule the most stable arrangement in a set of
degenerate orbitals is that with the most number of unpaired electrons. So
carbon has two unpaired electrons.
The electron configuration for carbon is 1s22s22p2. Notice the electron
configuration does not clearly indicate the number of unpaired electrons in
the element. The number of unpaired electrons is evident from the orbital
diagram. The orbital energy diagram for carbon is shown below.
Nitrogen has seven electrons. The placement of the next electron must
follow Hund's rule. The orbital diagram shows three unpaired electrons.
The electron configuration for nitrogen is 1s22s22p3.
For oxygen the eighth electron must pair with one of the electrons in the 2p
orbitals. The orbital diagram for oxygen is shown on the left. The electron
configuration for oxygen is 1s22s22p4.
The orbital diagrams for fluorine and neon are shown. The next two
electrons continue to pair those electrons that are unpaired to fill up the 2p
orbitals.
With neon the second level is filled with electrons. Completed levels are a
characteristic of all noble gases. If we look at the energy level diagram for
neon the completed second level means the next electron must go into the
third level. In the hydrogen atom the three sublevels, 3s, 3p and 3d were all
degenerate in energy. In the multi-electron atom the three sublevels do not
have the same energy. The relative energies of the three sublevels again
depend on the ability of the electron to penetrate to the nucleus. As in the
case of the second level the 3s orbital is lower in energy than the 3p which
is lower in energy compared to the 3d.
So as we progress from sodium across the period to argon the electrons are
placed in the orbitals just as they were for the second period.
The orbital diagrams for the eight elements are shown below.
The electron configurations for the next eight elements are as follows:
Na 1s22s22p63s1
Mg 1s22s22p63s2
Al 1s22s22p63s23p1
Si 1s22s22p63s23p2
P 1s22s22p63s23p3
S 1s22s22p63s23p4
Cl 1s22s22p63s23p5
Ar 1s22s22p63s23p6
When we come to potassium more interesting changes are observed.
Chemically potassium behaves like sodium, as an alkali metal. It appears
the next electron is in an s orbital, not a 'd' orbital. It turns out the energy
of the 4s orbital is very close to the energy of the 3d orbital at potassium.
But the energy of the 4s orbital is lower in energy compared to the 3d. So
the next electron is placed into the 4s orbital. At calcium the electron is
paired. For scandium we might consider whether the electron goes into the
3d or the 4p. It turns out the energy of the 3d is lower than the 4p so the d
sublevel begins to fill with scandium. The electron configuration for
scandium is [Ar]4s23d1. As electrons are added in titanium and vanadium
the configuration is [Ar]4s23d2 and [Ar]4s23d3. The next element,
chromium, would be expected to have a configuration of [Ar]4s23d4,
however this is not the case. It turns out that as a result of the similarity in
energy of the 4s sublevel and the 3d sublevel in this group that an
interesting phenomena occurs at chromium. Instead of [Ar]4s23d4 the
electron configuration is [Ar]4s13d5. We might suggest that a half-filled d
sublevel has extra stability. The next element, manganese, the additional
electron is added to complete the half-filled 4s sublevel and the
configuration is [Ar]4s23d5. From iron through nickel ([Ar]4s23d6,
[Ar]4s23d7, [Ar]4s23d8) the electrons spin-pair in the 3d sublevel. At copper
another reversal occurs. The electron configuration for copper is
[Ar]4s13d10 not [Ar]4s23d9. Zinc has an electron configuration of
[Ar]4s23d10.
At gallium we begin filling the 4p sublevel and continue to krypton.
Rubidium fills the 5s, yttrium the 4d and indium the 5p. Cesium fills the 6s
and lanthanum bigins the first available f sublevel, the 4f. The f sublevel is
filled from lanthanum through ytterbium. Throughout this period there
are strange reversals of configurations. The details of these changes are not
critically important to us.
Having gone through this exercise it is interesting to study the periodic
table in light of the position of the valance electrons of each atom. The
periodic table displayed uses color to denote the location of the outer-most
electrons. All of the alkali metals have the outer most, or valence electron,
in an s orbital. For lithium the outer most electron is in the 2s orbital, for
sodium the 3s, for potassium the 4s, etc. We say the general electron
configuration for the alkali metals is ns1. For the alkaline earth elements it
is ns2.
The periodic table can be used to write the electron configuration for any
element. The trick is locate the particular element in the correct level and
sublevel. The level numbers are located to the left of each period. The
sublevels are identified by noting the section of the periodic table the
element is located. (Describe the colors of the periodic table and do an
example using the periodic table.)
It is useful to understand the observed trends in particular physical
properties of the elements in relation to their location in the periodic table.
The idea is note the physical properties and then understand the observe
behavior in terms of the nature of the electron.