IN THIS ARTICLE YOU WILL LEARN ABOUT . . .
History and Arrangement of the Periodic Table
Metals, Nonmetals, and Semimetals
HISTORY OF THE PERIODIC TABLE
The periodic table has been developed and perfected over many years. Al-
though there are many scientists who have contributed to the periodic table,
the two scientists who are given the most credit are Dmitry Mendeleyev and
Henry Moseley. Mendeleyev, even though his periodic table had elements
missing from it, is given the most credit for the periodic table and periodic
trends. Later on, Moseley used a technique called x-ray crystallography and
discovered the idea of the atomic number. This discovery is the basis for the
arrangement of the modern periodic table. There were proposed periodic ta-
bles based upon atomic mass, but these arrangements did not sufﬁce because
of isotopes that can exist for an element.
ARRANGEMENT OF THE PERIODIC TABLE
The periodic table contains a number of periods and groups. The periods are
the horizontal rows. They are numbered 1 through 7. The groups (or fami-
lies) are the vertical columns. They are numbered 1 through 18. You will be
provided with a periodic table when you take the SAT II: Chemistry test.
METALS, NONMETALS, AND SEMIMETALS
Two categories of elements on the periodic table are the metals and the non-
metals. Their properties are summarized in the chart below:
• Are ductile and can be rolled into
• Are malleable and can be hammered
into thin sheets
• Conduct heat
• Conduct electricity
• Have a shiny luster
• Tend to lose electrons and become cations
• Make up two-thirds of the periodic table
• Are soft and brittle
• Lack luster
• Are poor conductors of heat
• Are poor conductors of electricity
• Tend to gain electrons and form
The semimetals, or metalloids, are known to exhibit some of the proper-
ties of metals and some of those of nonmetals. The semimetals are B, Si, Ge,
As, Sb, Te, and At. They are highlighted in bold in the partial periodic table
in Figure 4.1. The elements located to the left of the semimetals are the met-
als; those to the right of the semimetals are the nonmetals. Identifying an
element as a metal, nonmetal, or semimetal is important in identifying peri-
odic trends and in identifying the types of bonds that atoms will form with
PROBLEM: Identify the following elements as metals, nonmetals, or semi-
metals: potassium, calcium, bromine, hydrogen, and neon.
Solution: K and Ca are located on the left side of the periodic table and
are metals. Br and Ne are on the right side of the semimetals and are
nonmetals. Hydrogen, although on the left side of the periodic table, is
a nonmetal. If you’re still not convinced about hydrogen, ask yourself
about the properties of hydrogen gas and see where those properties ﬁt
in the comparison chart on p. 77.
Some groups or families are given special names and have certain properties
that should be addressed. But ﬁrst you must understand why elements are
put into the same group. Think about a family you know, not a chemical fam-
ily, but a human family. Children look like their parents. They learn to do
things from their parents and do them in the same way. The same holds true
for the elements in the families of the periodic table; they react the same way
(for the most part). As you learned in the last chapter, each element has a cer-
tain number of valence electrons. As you will learn in the next chapter, it is
the number of valence electrons of an atom that determines its chemical re-
activity. Because the elements in a family have the same number of valence
electrons, they will have a similar chemical reactivity. For example, Na and
K can be compared in electron conﬁguration and ions formed:
Both atoms have 1 valence electron and will lose this one electron to form
ions with charges of 1+. This similar charge will mean that both elements
have a similar chemical reactivity.
The important families and groups are listed below followed by their im-
portant characteristics. These characteristics will become more familiar to
you as you study the chapter on bonding.
Alkali Metals : Group 1
All group 1 metals have one valence
electron. When they form ions, they will
have a charge of 1+. Group 1 alkali metals
are highly reactive and will react vigor-
ously with water.
Alkaline Group 2 Earth Metals
All group 2 metals have two valence elec-
trons. When they form ions, they will have
a charge of 2+. Group 2 alkaline earth met-
als are highly reactive and will react with
Transition Groups 3–10, Metals d block
Transition metals are famous for the col-
ored salts and colored solutions they form.
Many gems contain numerous transition
metals. It is hard to predict the charge of a
transition metal ion because the transition
metals have multiple oxidation states. One
transition metal, Hg, exists as a liquid at
Halogens Group 17 Halogens (salt formers) have seven valence
electrons and form ions with a charge of
1−. The halogens exist in three phases at
room temperature. Fluorine is a pale-
yellow gas, chlorine is a green gas,
bromine is a brown-orange liquid, and
iodine is a purple solid.
Noble (Inert) Group 18 Gases
Noble gases have a full outer shell and will
not react to form ions or share electrons.
Lanthanides f Block and Actinides
These elements have their valence elec-
trons located in the f orbitals and are
radioactive in nature.
There are important periodic trends that occur across the periods and up
and down the groups. It is best to remember the trends of just a few elements.
This will simplify the trends greatly and make the periodic trend questions
the easiest to answer on the test.
Electronegativity is a measure of an atom’s ability to attract electrons. The
electronegativities of the elements are given a value of between 0.0 and 4.0.
The greatest electronegativity value goes to ﬂuorine, 4.0. So where is the ele-
ment with the lowest electronegativity? Look furthest from ﬂuorine and
across to the bottom left of the periodic table. Francium, Fr, has an electro-
negativity of 0.7. This should make sense because nonmetals tend to gain
electrons and have a higher electronegativity value, whereas metals tend to
lose electrons and have a lower electronegativity value. Because they don’t
react, the noble gases do not have a value for electronegativity.
PROBLEM:Which is expected to have a lower electronegativity, Na or S?
Solution: Na has a lower electronegativity because it is further from ﬂu-
orine on the periodic table.
Ionization energy, as its name suggests, is the energy needed to remove an
electron from an atom and form an ion. This concept should be easy to rec-
ognize in the periodic table once you have grasped the idea of electronega-
tivity. It takes a lot of energy to remove electrons from the very stable octets
of the noble gases. For example, for helium the ﬁrst ionization energy is 2372
kJ/mol, whereas neon has a ﬁrst ionization energy of 2081 kJ/mol. Fluorine,
with the highest electronegativity and the ability to “hold onto” electrons, has
a ﬁrst ionization energy of 1681 kJ/mol.
You might have guessed by now that the opposite holds true for the met-
als as you move further away from ﬂuorine and the noble gases. The proof
lies in the first ionization energies for iron (762 kJ/mol) and potassium
(419 kJ/mol). These values are just a fraction of the ﬁrst ionization energies
for certain nonmetals.
PROBLEM: Which is expected to have a greater ionization energy, Ca or Br?
Solution: Br is located closer to F and will have a higher ionization energy.
The atomic radius of an atom can be deﬁned as the distance from its nucleus
to the outermost electron of that atom. As you go down a group, the radius
of the atoms will increase as the atoms ﬁll more principal energy levels with
electrons. The proof for this trend can be seen in lithium, which has an
atomic radius of 155 picometers (10^−12 meters), and cesium, which has an
atomic radius of 267 picometers. You might expect the same to happen as
you examine the elements from left to right across a period. If lithium has
fewer electrons than ﬂuorine, then lithium should have a smaller radius than
ﬂuorine, right? Wrong! Fluorine has nine electrons and lithium has just
three, yet fluorine has an atomic radius of 57 picometers and lithium a
radius 155 picometers. Why the difference? Fluorine has more protons and
positive charge in its nucleus than does lithium. It turns out that when look-
ing at atomic radii across a period, it is the nuclear charge (and not the num-
ber of electrons) that determines the radius of the atom.
As covered in the previous chapter, atoms can gain or lose electrons. The
resulting ions can be expected to be of a different radius than that of the orig-
inal atom. When a nonmetal gains an electron, the ionic radius of the anion
will be bigger than that of the nonmetal atom. This is shown in Figure 4.2.
The opposite holds true for metal atoms and cations. Metals lose electrons
and will experience a decrease in their radius as shown in Figure 4.3.
THE s, p, d, AND f BLOCKS
The location of an element on the periodic table can tell a lot about the num-
ber of valence electrons the element has and in which subshell these valence
electrons can be located. These blocks are outlined in Figure 4.4.
The alkali and alkaline earth metals have their valence electrons in the s
subshells. Groups 13 through 18 have their valence electrons located in the
p subshells. The transition elements have their valence electrons in the d sub-
shells, and ﬁnally, the lanthanides and actinides have their valence electrons
in the f sublevel.
1. The modern periodic table is arranged
based upon atomic
2. In period 3 of the periodic table the atom
with the largest atomic radius is located in
3. The elements that display the greatest non-
metallic character are located toward which
corner of the periodic table?
(A) Upper left
(B) Dead center
(C) Lower right
(D) Lower left
(E) Upper right
4. Which two elements will display the most
similar chemical properties?
(A) Aluminum and calcium
(B) Nickel and phosphorus
(C) Chlorine and sulfur
(D) Carbon and sulfur
(E) Lithium and potassium
5. Assuming the ground state, all of the ele-
ments located in group 13 of the periodic
table will have the same number of
(A) nuclear particles
(B) occupied principal energy levels
(D) valence electrons
6. Which group contains elements in the solid,
liquid, and gas phases at 298 K and 1 atm?
7. An element that has a high ﬁrst ionization
energy and is chemically inactive would
most likely be
(A) a noble gas
(B) a transition element
(C) an alkali metal
(D) a halogen
(E) an alkaline earth metal
8. Which salt solution is most likely to be col-
(A) KClO3 (aq)
(B) KNO3 (aq)
(C) K2CrO4 (aq)
(D) K2SO4 (aq)
(E) KCl (aq)
9. As the elements of period 2 are considered
from left to right, there is generally a de-
(A) ionization energy
(C) metallic character
(D) nonmetallic character
(E) none of the above
10. Which element is a liquid at room tempera-
11. At STP, which element is most expected to
exist as a monatomic gas?
12. Nonmetals are poor conductors of heat and
they also tend to
(A) be brittle
(B) conduct an electrical current
(C) have a shiny luster
(D) be malleable
(E) lose electrons
13.Which statement does not explain why
elements in a group are placed together?
(A) They tend to have the same number of
(B) They tend to have a similar oxidation
(C) They tend to have the same electroneg-
(D) They tend to have the same chemical
(E) They tend to have the same charge when
they form ions.
1. (B) 2. (A) 3. (E)
4. (E) 5. (D) 6. (D)
7. (A) 8. (C) 9. (C)
10. (B) 11. (D) 12. (A)