The Periodic Table and Periodic Properties
We just saw how the periodic table can help us quickly
determine electron configurations and quantum numbers. As you’ll
see in this section, this is possible because of the special arrangement
of elements in the periodic table. There will almost definitely
be at least one question about trends in the periodic table on the
SAT Chemistry test, so be sure to read this section closely.
The Anatomy of the Periodic Table
As you are probably well aware, in the periodic table,
elements are arranged in order of increasing atomic number. The
18 vertical columns of the table are called groups or families,
while the seven horizontal rows are called periods and
correspond to the seven principal quantum energy levels, n =
1 through n = 7.
On the right side of the periodic table is a dividing
line resembling a staircase. To the left of the staircase lie the
metals, and to the right of the staircase lie the nonmetals. Many
of the elements that touch the staircase are called metalloids,
and these exhibit both metallic and nonmetallic properties. Study
the diagram below and memorize the names of the different types
of elements, because you will definitely see questions about these
groupings on the test!
Metals are malleable, ductile, and have luster;
most of the elements on the periodic table are metals. They oxidize
(rust and tarnish) readily and form positive ions
(cations). They are excellent conductors of both heat and electricity.
The metals can be broken down into several groups.
Transition metals (also called the transition
elements) are known for their ability to refract light as a result
of their unpaired electrons. They also have several possible oxidation
states. Ionic solutions of these metals are usually colored, so
these metals are often used in pigments. The actinides and lanthanides
are collectively called the rare earth elements and
are filling the f orbitals. They are rarely found
in nature. Uranium is the last naturally occurring element; the
rest are man-made.
Nonmetals lie to the right of the staircase
and do not conduct electricity well because they do not have free
electrons. All the elemental gases are included in the nonmetals. Notice
that hydrogen is placed with the metals because it has only one
valence electron, but it is a nonmetal.
Here are some specific families you should know about,
within the three main groups (metals, nonmetals, and metalloids):
Alkali metals (1A)—The most
reactive metal family, these must be stored under oil because they
react violently with water! They dissolve and create an alkaline,
or basic, solution, hence their name.
Alkaline earth metals (2A)—These also are
reactive metals, but they don’t explode in water; pastes of these
are used in batteries.
Halogens (7A)—Known as the “salt formers,”
they are used in modern lighting and always exist as diatomic molecules
in their elemental form.
Noble gases (8A)—Known for their extremely
slow reactivity, these were once thought to never react; neon, one
of the noble gases, is used to make bright signs.
Now that you’re familiar with the different groupings
of the periodic table, it’s time to talk about the ways we can use
the periodic table to predict certain characteristics of elements.
Since in an atom there is no clear boundary beyond which
the electron never strays, the way atomic radius is measured is
by calculating the distance between the two nuclei of atoms when
they are involved in a chemical bond. If the two bonded atoms are
of the same element, you can divide the distance by 2 to get the
atom’s radius. That said, one of the two important things you’ll
need to know about atomic radii for the SAT II Chemistry exam is that atomic
across a period from left to right
. But why? It seems as though
the more protons you add, the more space the atom should take up,
but this is not the case. The reason for this lies in the basic
concept that opposite charges attract each other and like charges
repel each other. As you increase the number of protons in the nucleus
of the atom, you increase the effective nuclear charge
the atom (Zeff
), and the nucleus
pulls more strongly on the entire electron cloud. This makes the
atomic radius decrease in size. The second thing you’ll need to
know is that atomic radii increase moving down a group or
. This is easier to understand if you refer to the
Bohr model. As you move down the table, the value of n
as we add another shell. Remember that the principal quantum number, n
determines the size of the atom. As we move down a family, the attractive
force of the nucleus dissipates as the electrons spend more time
farther from the nucleus.
One more thing about atomic size. As you know, when an
atom loses an electron, a cation, or positive ion,
is formed. When we compare the neutral atomic radius to the cationic radius,
we see that the cationic radius is smaller. Why? The protons in
the nucleus hold the remaining electrons more strongly. As you might
expect, for negatively charged ions, or anions, the
nuclear attractive force decreases (and there is enhanced electron-electron repulsion),
so the electrons are less tightly held by the nucleus. The result
is that the anion has a larger radius than the neutral atom.
The SAT II Chemistry test might ask you to compare the
sizes of two atoms that are isoelectronic, meaning
that they have the same number of electrons. In this case, you would then
consider the number of protons the two atoms possess.
Which ion is larger, F– or
Since these two atoms are isoelectronic and in the same
period, the atom with more protons in its nucleus will hold its
electrons more tightly and be smaller. Fluoride will be smaller
since it has more protons (9, compared to oxide’s 8).
Ionization Energy (IE)
The ionization energy of an atom is the energy
required to remove an electron from the atom in the gas phase. Although
removing the first electron from an atom requires energy, the removal
of each subsequent electron requires even more energy. This means
that the second IE is usually greater than the first, the third
IE is greater than the second, and so on. The reason it becomes
more difficult to remove additional electrons is that they’re closer
to the nucleus and thus held more strongly by the positive charge
of the protons.
Ionization energies differ significantly, depending on
the shell from which the electron is taken. For instance, it takes
less energy to remove a p electron than an s electron,
even less energy to extract a d electron, and the
least energy to extract an f electron. As you can probably
guess, this is because s electrons are held closer
to the nucleus, while f electrons are far from
the nucleus and less tightly held. You’ll need to remember two important
facts about ionization energy for the test. The first is that ionization
energy increases as we move across a period.
The reason for this, as is the case with periodic trends
in atomic radii, is that as the nucleus becomes more positive, the
effective nuclear charge increases its pull on the electrons and
it becomes more difficult to remove an electron.
The second thing you’ll need to remember is that ionization
energy decreases as you move down a group or family. The
increased distance between electrons and the nucleus and increased
shielding by a full principal energy level means that it requires
less energy to remove an electron. Shielding occurs
when the inner electrons in an atom shield the outer electrons from
the full charge of the nucleus. Keep in mind that this phenomenon
is only important as you move down the periodic table! Here are
the values for the first ionization energies for some elements:
There are some important exceptions to the above two ionization
energy trends in the periodic table, so make sure you study these
- When electron pairing first occurs within
an orbital, electron-electron repulsions increase, so that removing
an electron takes less energy (it’s easier); thus the IE drops at
this time. For example, less energy is required to remove an electron
from oxygen’s valence in spite of an increasing Zeff because
oxygen’s p4 electron
is the first to pair within the orbital. The repulsion created lowers
the amount of energy required to remove either electron.
- There is also a drop in ionization energy from s2 to p1—also in
spite of an increasing Zeff.
This drop is due to the fact that you are removing a p electron rather
than an s electron. The p electrons
are less tightly held because they do not penetrate
the electron cloud toward the nucleus as well as an s electron
Which of the following elements has the highest ionization
energy: K, Ca, Ga, As, or Se?
The answer is arsenic, or As. Since IE increases as we
move across a period, you may have chosen Se. However, there is
a drop in IE in spite of increasing Zeff due
to the increased electron-electron repulsion in the family that
contains oxygen, since they are np4.
An atom’s electron affinity is the amount
of energy released when an electron is added to the atom in its
gaseous state—when an electron is added to an atom, the atom forms
a negative ion. Most often, energy is released as
an electron is added to an atom, and the greater the attraction
between the atom and the electron added, the more negative the atom’s
For the SAT II Chemistry test, remember that electron
affinity becomes more negative as we move across a period.
This means that it’s easier to add an electron to elements, the farther
to the right you travel on the periodic table. Why? Again, this
is because the higher Zeff increases
the nuclear attraction for the incoming electron. Important exceptions
to this rule are the noble gases: He, Ne, Ar, Kr, and Xe. They have
electron affinities that are positive (meaning very low), because
if they were to accept another electron, that electron would have
to go into a new, higher-energy subshell, and this is energetically
Electron affinities do not change very much as
you go down a group. This is because the lower electron-nucleus
attraction that’s seen as we go down a group is pretty evenly counterbalanced
by a simultaneous lowering in electron-electron repulsion. Remember
that there is no clear trend for electron affinity as you go down
a group on the periodic table—this fact could come up in a synthesis
of knowledge question!
Electronegativity is a measure of the attraction an atom
has for electrons when it is involved in a chemical bond. Elements
that have high ionization energy and high electron affinity will
also have high electronegativity since their nuclei strongly attract
electrons. Electronegativity increases from left to right as we
move across a period and decreases as we move down any group or
By now, these trends should make sense. You know that
ionization energies tend to decrease with increasing atomic number
in a group, although there isn’t a significant change in electron
affinity, so it makes sense that atoms’ attraction for electrons
in a bond would also increase as their Zeff increased.
We will discuss the concept of electronegativity further in the
next section, when we discuss chemical bonding.
Here’s a summary of the trends we discussed in this section.
Make sure to memorize them!