Periodic Trends Around the Table
Atomic Radius: Not as simple as defining the radius in geometry class because atomic orbitals are
not a fixed distance from the nucleus. Atomic radius still expresses the size of an atom (larger radius,
larger atom) but a better way to define it is as "half of the distance between the centers of two atoms
of the same element that are just touching each other."
Ionization Energy: Basically the minimal energy needed or "energy cost' to pull electrons off of
gaseous atoms. Since elements in Group 1A have 1 valence electron they want to get rid of anyway,
the ionization energy required for another atom to take that valence electron is minimal. Good luck
getting a valence electron from a Noble Gas on the other hand! They have the highest ionization
energy on the Periodic Table because they have the number of valence electrons they need to fill their
s and p orbitals. In general, ionization energies measure the tendency of a neutral atom to resist the
loss of electrons.
Electron Affinity: The yin to ionization energy's yang, electron affinity is the energy given off when
a neutral atom in the gas phase grabs an extra electron and forms a anion. That's right, gain
electron/s, give off energy.
Electronegativity: While ionization energy and electron affinity measure energy transfer to move
electrons around, electronegativity is (to personify atoms a bit) a "desire" scale. Electronegativity
measures an atom's tendency (desire) to attract boding electron/s. Fluorine's is the highest while
Francium's is the lowest.
Chemical Reactivity: Reactivity tells you how likely or vigorously atoms will react with each other.
Atomic radius, ionization energy (how easily electrons can be removed) and electronegativity (how
badly atoms want to steal other atom's electrons) are great predictors because electron transfers are
the basic for inorganic chemical reactions.
Chemical Reactivity Let’s consider how metals react. Metals react by losing electrons. They have a
low ionization energy so it's fairly easy for them to lose electrons.
Metal reactivity: As you go across a period, the nuclear charge will increase; the number of
energy levels will stay the same, so there is a stronger and stronger attraction for the electrons. The
electrons are being held more tightly as you go across a period. It becomes more and more difficult
to lose electrons and consequently the reactivity of the metals decreases as you go from left to right
across the periodic table.
As you go down the periodic table, the nuclear charge increases but so does the number of shielding
electrons. Consequently, the dominant factor is that we have more and more energy levels and the
electrons are further and further away from the nucleus. Thus it is easier for those electrons to come
off.
Nonmetal reactivity: this trend runs opposite of metal reactivity. Nonmetals react differently than
metals. They react by losing electrons.
Relating Reactivity of Nonmetals to Atomic Structure
As you go down a nonmetallic group in the periodic table, the elements become less reactive. You
also know that as you go down a group on the periodic table, the number of energy levels is the
most predominant factor. If an electron comes into an atom that has a large number of energy levels,
it will be further away from the nucleus and not be attracted as strongly as it would be in a smaller
atom with fewer energy levels. So the closer to the nucleus, the more the attraction is felt, the more
reactive that element is.
The reactivity of the nonmetals ties in well with the concept of electron affinity and the tendency to
gain electrons. With nonmetals, the greater the tendency to gain electrons, the more reactive it is.
This argument should hold true whether we are talking about nonmetals within a family or within a
period. As you go across a period, there is a greater nuclear charge and thus the electrons should
be attracted more readily by elements that are further to the right and the tendency to gain electrons
will increase. Thus the reactivity of the nonmetals should increase as you go from left to right across
the periodic table, up to but not including the inert gases.
Atomic Size
Let's make some comparisons in a family and in a period. As you go down a group, like from hydrogen to
lithium to sodium on down, the atomic size increases. As you go across a period, as from lithium to neon,
notice that the size decreases. You need to remember (or memorize) those trends.
Remember that the nuclear charge and the shielding electrons combine to make the effective nuclear
charge. That is a very important factor when you are comparing elements in a period. As you go across a
period, the nuclear charge increases and the number of energy levels stays the same. Consequently, the
number of shielding electrons stays the same and the effective nuclear charge increases. As the effective
nuclear charge increases, it pulls the electrons in closer and closer to the nucleus. So as you go across a
period, the increase in the nuclear charge causes a decrease in the atomic size because the electrons in the
valence energy level are pulled in closer and closer.
Now let's make comparison within a group such as hydrogen down to francium (Fr). It is true that the
nuclear charge is increasing, but so is the number of shielding electrons. The number of shielding
electrons increases by the same amount that the nuclear charge increases. So the effective nuclear
charge felt by the valence electrons stays the same. There is no increase in the effective nuclear
charge but there is an increase in the number of energy levels that are being used. Consequently,
the electrons in the valence energy level are farther and farther away from the nucleus because they
are in higher energy levels. Consequently, the important factor in a vertical comparison on the
periodic table is the number of energy levels that are being used because the increase in the number
of shielding electrons cancels out the increase in the nuclear charge.
To summarize, as you go across a period, the increase in the nuclear charge is the most important
factor because the number of energy levels stays the same. As you go down a group, the increase
in shielding electrons more or less cancels out the increase in nuclear charge, leaving the increase
in the number of energy levels as the most important factor. This is true not only for atomic size
but for other properties as well.
Practice with Comparing Atomic Size For each of the following sets of atoms, decide which is
larger, which is smaller, and why.
a. Li, C, F
b. Li, Na, K
c. Ge, P, O
d. C, N, Si
e. Al, Cl, Br
a. Li, C, F: All are in the same period and thus have the same number of energy levels. Therefore,
the important factor is the nuclear charge. Li is the largest because it has the smallest nuclear
charge and pulls the electrons toward the nucleus less than the others. F is the smallest because it
has the largest nuclear charge and pulls the electrons toward the nucleus more than the others.
b. Li, Na, K: All are in the same group and thus have the same effective nuclear charge. Therefore,
the important factor is the number of energy levels. Li is the smallest because it uses the smallest
number of electron energy levels. K is the largest because it uses the largest number of electron
energy levels.
c. Ge, P, O: All are in different groups and periods, therefore both factors must be taken into
account. Fortunately both factors reinforce one another. Ge is the largest because it uses the largest
number of energy levels and has the smallest effective nuclear charge. O is the smallest because it
uses the smallest number of energy levels and has the largest effective nuclear charge.
d. C, N, Si: Not all are in the same group and period, so, again, both factors must be taken into
account. C and N tie for using the smallest number of energy levels, but N has a higher effective
nuclear charge. Therefore, N is the smallest. C and Si tie for having the lowest effective nuclear
charge, but Si uses more energy levels. Therefore, Si is the largest.
e. Al, Cl, Br: Not all are in the same group and period, so, again, both factors must be taken into
account. Cl is the smallest because it has higher effective nuclear charge than Al and uses fewer
energy levels than Br. Which is largest is less straightforward. Al has a lower effective nuclear
charge (by four), but Br uses more energy levels (by one). Because the difference in effective
nuclear charge is larger, it should be the more important factor in this case, making Al the largest.
Al and Br can also be compared to one another indirectly by comparing both to Cl. Both Al and Br
are larger than Cl. Al is larger than Cl because it has lower effective nuclear charge (by four). Br is
larger than Cl because it uses more energy levels (by one). Because Al is larger than Cl by four
"steps" and Br is larger than Cl by only one "step", Al is likely the largest of the three.
Ionization Energy
If the ionization energy is high, that means it takes a lot of energy to remove the outermost
electron. If the ionization energy is low, that means it takes only a small amount of energy to
remove the outermost electron.
Let’s use your understanding of atomic structure to make some predictions. Think for a minute about
how ionization energy would be affected by three of the factors we were talking about earlier: (1)
nuclear charge, (2) number of energy levels, and (3) shielding.
As the nuclear charge increases, the attraction between the nucleus and the electrons increases
and it requires more energy to remove the outermost electron and that means there is a higher
ionization energy. As you go across the periodic table, nuclear charge is the most important
consideration. So, going across the periodic table, there should be an increase in ionization energy
because of the increasing nuclear charge.
Going down the table, the effect of increased nuclear charge is balanced by the effect of increased
shielding, and the number of energy levels becomes the predominant factor. With more energy
levels, the outermost electrons (the valence electrons) are further from the nucleus and are not so
strongly attracted to the nucleus. Thus the ionization energy of the elements decreases as you go
down the periodic table because it is easier to remove the electrons.
This table shows the measured values for the
ionization energies of the first twenty elements. If
you take a close look at what happens to the
H
ionization energy as you go from left to right
13.6
across the periodic table, you will find that there
is not really a steady increase in ionization
Li
energy as I had indicated. You could describe the 5.4
pattern you see there as being a few steps
Na
forward then one step back, repeating itself as
5.1
you move across. Progress is made, but it is not
steady.
Ionization Energies (v)
He
24.6
Be
9.3
B
8.3
C
11.3
N
14.5
O
13.6
F
17.4
Ne
21.6
Mg
7.6
Al
6.0
Si
8.2
P
10.5
S
10.4
Cl
13.0
Ar
15.8
The periodic nature of ionization energy is emphasized in this diagram. With each new period the
ionization energy starts with a low value. Within each period you will notice that the pattern is really
kind of a zigzag pattern progressing up as you go across the periodic table.
The zigs and zags on that graph
correspond to the sublevels in the
energy levels. So far in this lesson we
have presumed that all the electrons in
the second energy level are pretty
much the same. Two factors make that
not completely true. One factor is that
because s and p orbitals have different
shapes, the electrons in p orbitals
have more energy and are further from
the nucleus. The other factor is that
when electrons are paired up in an
orbital, they repel one another
somewhat. Those two factors account
for the zigzag nature of the increase in
ionization energy.
Nevertheless, as a general trend, from left to right across the periodic table, ionization energy does
increase. Also as you go down the periodic table, the ionization energy does decrease for the
reasons given.
Practice with Comparing Ionization Energies
For each of the following sets of atoms, decide which has the highest and lowest ionization energies
and why.
a.
b.
c.
d.
e.
Mg, Si, S
Mg, Ca, Ba
F, Cl, Br
Ba, Cu, Ne
Si, P, N
Answers to Comparing Ionization Energies
a. Mg, Si, S: All are in the same period and use the same number of energy levels. Mg has the
lowest I.E. because it has the lowest effective nuclear charge. S has the highest I.E. because it has
the highest effective nuclear charge.
b. Mg, Ca, Ba: All are in the same group and have the same effective nuclear charge. Mg has the
highest I.E. because it uses the smallest number of energy levels. Ba has the lowest I.E. because it
uses the largest number of energy levels.
c. F, Cl, Br: All are in the same group and have the same effective nuclear charge. F has the
highest I.E. because it uses the smallest number of energy levels. Br has the lowest I.E. because it
uses the largest number of energy levels.
d. Ba, Cu, Ne: All are in different groups and periods, so both factors must be considered.
Fortunately both factors reinforce one another. Ba has the lowest I.E. because it has the lowest
effective nuclear charge and uses the highest number of energy levels. Ne has the highest I.E.
because it has the highest effective nuclear charge and uses the lowest number of energy levels.
e. Si, P, N: Si has the lowest I.E. because it has the lowest effective nuclear charge and is tied (with
P) for using the most energy levels. N has the highest I.E. because it uses the fewest energy levels
and is tied (with P) for having the highest effective nuclear charge.
Electron Affinity Tendency to Gain Electrons
Next let's consider the opposite of losing electrons, which is, of course, the gaining of electrons. Atoms can
attract additional electrons if there is room for them in the valence energy level. When an extra electron
moves into the valence shell, it can feel the attraction exerted by the effective nuclear charge. Because the
effective nuclear charge is largest for the elements on the right side of the periodic table, those atoms
provide the greatest attraction for electrons and have the greatest tendency to gain electrons.
Thus the tendency of atoms to gain electrons increases as we go from left to right across the periodic table.
At least it increases until we get to the inert gases. There it drops off to zero because there is no room for
additional electrons in the valence energy level. A new electron would have to start a new energy level, but
there would not be an additional proton in the nucleus to provide any effective nuclear charge.
As we look at elements going down the periodic table, the effective nuclear charge remains the same, so the
increase in the number of energy levels is the important factor. The tendency of atoms to gain electrons
decreases as we go down the periodic table. The reason for this is simply that with the larger atoms the
added electron is not as close to the nucleus and therefore the attractive force exerted by the effective nuclear
charge is not as powerful as it is in the smaller atoms.
Practice Comparing Tendencies to Gain Electrons
For each of the following sets of atoms, decide which has the least and which has the greatest tendency to
gain electrons and why. Check your answers below and then continue with the lesson.
a. Li, C, N
b. C, O, Ne
c. Si, P, O
d. K, Mg, P
e. S, F, He
Answers for Comparing Tendencies to Gain Electrons
a. Li, C, N: Li has the least tendency to gain electrons because it has the lowest effective nuclear charge (and
all use the same number of energy levels). N has the greatest tendency to gain electrons because it has the
highest effective nuclear charge (and all use the same number of energy levels).
b. C, O, Ne: Ne has the lowest tendency to gain electrons because its outer energy level is full and there is
no room for an additional electron. O has the greatest tendency to gain electrons because it has a higher
effective nuclear charge than C (and both use the same number of energy levels).
c. Si, P, O: O has the greatest tendency to gain electrons because it has the highest effective nuclear charge
and uses the smallest number of energy levels. Si has the lowest tendency to gain electrons because it has
the lowest effective nuclear charge and is tied (with P) for using the most energy levels.
d. K, Mg, P: P has the greatest tendency to gain electrons because it has the highest effective nuclear charge
and is tied (with Mg) for using the smallest number of energy levels. Neither Mg nor K have much attraction
for electrons, but K has the lowest tendency to gain electrons because it has the lowest effective nuclear
charge and uses the most energy levels.
e. S, F, He: He has the lowest tendency to gain electrons because its outer energy level is full and there is no
room for an additional electron. F has the greatest tendency to gain electrons because it has a higher
effective nuclear charge and uses fewer energy levels than S.