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IMPORTANT ASPECTS OF ATOMIC STRUCTURE
| Next let's consider several important aspects of atomic structure that
influence the periodic atomic properties and the chemical properties of the different
elements. (These are also listed for you in objective 7.) |
Nuclear Charge
Main Electron Energy Levels
Shielding Electrons
Valence Electrons
Effective Nuclear Charge |
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NUCLEAR CHARGE
The first of these, and the simplest, is the nuclear charge. The nuclear charge
is the total charge of all the protons in the nucleus. It has the same value as the atomic
number. The nuclear charge increases you go across the periodic table. As you get to the
end of one period and you go to the beginning of the next period, the atomic number and
the nuclear charge continues to increase. Consequently, there is no periodic or repeating
nature to the nuclear charge. The nuclear charge just keeps increasing. This is true
whether you go across the period or down a group.
MAIN ENERGY LEVELS
The number of main energy levels for electrons is another very important
consideration. As the atomic number increases, so does the number of electrons. But that
does not necessarily increase the number of energy levels.
As you go across a period all of the new electrons fit into
the energy levels that are already being used. For example, looking at carbon, nitrogen,
and oxygen, the number of protons increases from 6 to 7 to 8; so does the number of
electrons. However, when we look at their electron arrangements, notice that all of the
electrons are in the first two energy levels. The number of energy levels being used does
not change even though the number of electrons does.
|
| |
C |
N |
O |
| electrons |
6 |
7 |
8 |
| configuration |
1s22s22p2 |
1s22s22p3 |
1s22s22p4 |
| levels |
2 |
2 |
2 |
|
| (This is also shown in example 5a in your workbook. The
format there shows the symbol for the element, the number of protons in the nucleus, the
electron configuration, and also the total number of electrons in each energy level or
shell.) |
It is only when you go from one period to the next that you have to
increase the number of energy levels. (Also shown in a different way in part b of example
5 in your workbook.). As we go from fluorine to neon to sodium, the number of protons
increases from 9 to 10 to 11 and thus the number of electrons increases from 9 to 10 to
11. Notice what happens to the number of energy levels that must be used. For both
fluorine and neon, two energy levels accommodate all of the electrons. But once there are
10 electrons in those two energy levels (2 in the first and 8 in the second as with neon),
any additional electrons have to go into the next energy level.
|
| |
F |
Ne |
Na |
| electrons |
9 |
10 |
11 |
| configuration |
1s22s22p5 |
1s22s22p6 |
1s22s22p63s1 |
| levels |
2 |
2 |
3 |
|
| Next, let’s consider what happens within a group. (Also shown in
example 5c in your workbook.) As you go from carbon to silicon to germanium, the number of
protons increases in large jumps. The number of electrons also increase and the number of
energy levels used also increases. Notice that carbon has two energy levels, silicon has
three, and germanium (Ge) has four levels being used. |
| |
electrons |
configuration |
levels |
| C |
6 |
1s22s22p2 |
2 |
| Si |
14 |
1s22s22p63s23p2 |
3 |
| Ge |
32 |
1s22s22p63s23p64s23d104p2 |
4 |
|
| The number of energy levels used to accommodate the electrons
in the atoms of a particular element is going to be the same as the number of the period.
As you go down the periodic table, you will increase the number of energy levels being
used. |
Like the atomic number, the number of energy levels is not really a periodic or
repeating feature of atoms. It stays the same throughout a period and then increases by
one when a new period starts.
VALENCE ELECTRONS
The valence electrons are the electrons in the last shell or energy level of an
atom. They do show a repeating or periodic pattern. The valence electrons increase in
number as you go across a period. Then when you start the new period, the number drops
back down to one and starts increasing again.
| For example, when you go across the table from carbon to nitrogen to
oxygen, the number of valence electrons increases from 4 to 5 to 6. As we go from fluorine
to neon to sodium, the number of valence electrons increases from 7 to 8 and then drops
down to 1 when we start the new period with sodium. Within a group--starting with carbon
and going down to silicon and germanium--the number of valence electrons stays the same. |
| |
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C
4 |
N
5 |
O
6 |
F
7 |
Ne
8 |
Na
1 |
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Si
4 |
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| |
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Ge
4 |
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So generally speaking, the number of valence electrons stays the same as you go up or
down a group, but they increase as you go from left to right across the periodic table.
The preceding statement works very well for the representative elements, but it comes a
bit short of the truth when you start talking about the transition elements.
Electrons going into the d sublevels of the transition metals complicate this pattern.
In some ways these electrons behave like valence electrons. In some other ways they behave
like shielding electrons, which are discussed in the next section. The first electrons
into a d sublevel seem to behave more like valence electrons but the last ones seem to act
more like shielding electrons, with variations along the way. Switching the order from
4s3d to 3d4s is one way to represent this.
| |
Sc |
Ti |
V |
Cr |
Mn |
Fe |
Co |
Ni |
Cu |
Zn |
| outer configuration |
4s23d1 |
4s23d2 |
4s23d3 |
4s13d5 |
4s23d5 |
4s2
3d6 |
4s2
3d7 |
4s2
3d8 |
3d104s1 |
3d104s2 |
| apparent valence elctrons |
3 |
2-4 |
2-5 |
2-6 |
2-7 |
2 or 3 |
2 or 3 |
2 or 3 |
1 or 2 |
2 |
SHIELDING ELECTRONS
Shielding electrons are the electrons in the energy levels between the
nucleus and the valence electrons. They are called "shielding" electrons because
they "shield" the valence electrons from the force of attraction exerted by the
positive charge in the nucleus.
| In fluorine there are 9 protons in the nucleus and there are 2 shielding
electrons in the first level between the nucleus and the outer shell. They shield some of
the charge of the nucleus from the electrons that are in the outermost energy level. (Also
look at example 5b in your workbook.) |
| |
nuclear
charge |
shielding
electrons |
valence
electrons |
| F |
+9 |
1s2 |
2s22p5 |
| 2 |
7 |
|
| Next, neon also has 2 shielding electrons along with 8 valence electrons. |
| |
nuclear
charge |
shielding
electrons |
valence
electrons |
| Ne |
+10 |
1s2 |
2s22p6 |
| 2 |
8 |
|
| With sodium, we have 3 energy levels. There is one valence electron in the
third level and all the electrons between that one valence electron and the nucleus are
shielding electrons. In this case there are 2 in the first energy level and 8 in the
second for a total of 10 shielding electrons. |
| |
nuclear
charge |
shielding
electrons |
valence
electrons |
| Na |
+11 |
1s22s22p6 |
3s1 |
| 10 |
1 |
|
| So notice that the number of shielding electrons increases when you reach
the end of the periodic table and go on to the next period. |
|
| Now look at carbon, nitrogen, and oxygen to see that within a period there
is no change in the number of shielding electrons. Even though the valence electrons
increase in number from 4 to 5 to 6, the number of shielding electrons stays the same--two
shielding electrons for each of those elements. |
| |
C |
N |
O |
| shielding electrons |
2 |
2 |
2 |
| valence electrons |
4 |
5 |
6 |
|
| When you deal with the changes within a group, notice what happens. (Also
see part c of example 5.) Going from carbon to silicon to germanium, the number of protons
in the nucleus increases from 6 to 14 to 32, the number of energy levels increases from 2
to 3 to 4, the number of shielding electrons also increases. In carbon there are 4 valence
electrons and 2 shielding electrons. Silicon also has 4 valence electrons, but it has 10
shielding electrons. Germanium (Ge) also has 4 valence electrons, and it has 3 shells or
energy levels of electrons that are shielding electrons. There are 2 in the first, 8 in
the second, and 18 in the third for a total of 28 shielding electrons along with the 4
valence electrons. |
| |
shielding
electrons |
valence
electrons |
| C |
2 |
4 |
| Si |
10 |
4 |
| Ge |
28 |
4 |
|
Notice that the shielding electrons follow a pattern somewhat like the number of energy
levels. They stay the same within a period (except for increasing gradually and
erratically across the transition metals). They increase in steps as you start a new
period or go down a group
EFFECTIVE NUCLEAR CHARGE.
The next thing to be considered is effective nuclear charge. Generally speaking,
effective nuclear charge is the charge felt by the valence electrons after you have taken
into account the number of shielding electrons that surround the nucleus.
| Again let's take a look at a fluorine atom. (Also note example 5b in your
workbook.) The nucleus itself has a +9 charge and anything in its vicinity will feel that
charge. The two electrons in the first energy level as they look at the nucleus feel a +9
charge because that is the charge on the nucleus. But the electrons that are in the
valence energy level would be shielded from the nucleus by the 2 shielding electrons. The
+9 nuclear charge is shielded by 2 electrons to give an effective nuclear charge of +7
that is felt by the valence electrons. If you get out beyond the valence electrons, then
the effective charge is 0 simply because the +9 charge of the nucleus is surrounded by 9
electrons. |
| |
nuclear charge |
shielding electrons |
valence electrons |
| F |
+9 |
1s2 |
2s22p5 |
| 2 |
7 |
| |
+7 |
7 |
| effective nuclear charge |
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Generally we are only concerned with the effective nuclear charge as it pertains to the
valence electrons (+7 in this case), but sometimes the broader concept of what charge is
felt by other electrons is useful.
QUICK QUIZ ON EFFECTIVE NUCLEAR CHARGE
With that in mind, figure out what the effective nuclear charge would be for neon and
sodium for the electrons in each energy level. What would the first two electrons feel?
What would the next eight electrons feel? Then for sodium, what would that last electron
feel? So take a moment to figure that out.
***What charge is felt by the electrons in the first level of a neon atom?
etc.***
EFFECTIVE NUCLEAR CHARGE CONTINUED
For neon you should have figured that the electrons in the first energy level would
feel a nuclear charge of +10, the full nuclear charge. Then the 8 valence electrons would
feel a +8 effective nuclear charge. Going on down to sodium, the first 2 electrons would
feel a +11 nuclear charge. The 8 electrons in the second energy level would feel a +9
effective nuclear charge. The one valence electron would feel a +1 effective nuclear
charge.
| Notice what happens. As you go across the table from fluorine to
neon, the effective nuclear charge felt by the valence electrons increases. Then as you go
to sodium in the next period, there is another energy level, the number of shielding
electrons increases, causing the effective nuclear charge felt by the valence electron to
drop. |
| Effective Nuclear Charge for Valence Electrons |
| F |
Ne |
Na |
| +7 |
+8 |
+1 |
|
Notice that in all these examples the effective nuclear charge is the same as the
number of valence electrons. That is true as long as you are dealing with neutral atoms.
However many valence electrons there are, that will be the effective nuclear charge that
the valence electrons feel. It has to be that way for neutral atoms. It is not true when
dealing with ions.
Also notice that the effective nuclear charge depends on both the nuclear charge and
the number of shielding electrons. The nuclear charge keeps increasing. Meanwhile, the
shielding electrons stay constant while you are going across s and p parts of the period,
(but increase gradually across the d part of the period). Then when you go to the next
period, they jump in number. Consequently, the effective nuclear charge drops at that
point. Therefore, the effective nuclear charge increases as you go across a period and
then drops and starts over again at +1 when you start the next period. Within a period the
effective nuclear charge increases as you go across the periodic table.
| As you go down a group, the increase in the nuclear charge is cancelled
out by the increase in shielding electrons and the effective nuclear charge stays pretty
much the same. In carbon the 4 valence electrons in the outermost shell feel a +6 charge
surrounded by two shielding electrons for a +4 effective nuclear charge. For silicon it
would also be a +4 effective nuclear charge because the 14 protons in the nucleus are
surrounded by 10 shielding electrons. Germanium (Ge) has 32 protons and it has 28
shielding electrons and so the valence electrons feel an effective nuclear charge of +4.
As you go down a group, the increase in the nuclear charge is balanced by an increase in
the number of shielding electrons so that the effective nuclear charge remains the same. |
| |
nuclear charge |
shielding electrons |
effective nuclear charge |
| C |
+6 |
2 |
+4 |
| Si |
+14 |
10 |
+4 |
| Ge |
+32 |
28 |
+4 |
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One other thing I should mention about the effective nuclear charge is that it is quite
often referred to as the kernel charge. The "kernel" includes the nucleus
and all shielding electrons but does not include the valence electrons.
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