Chapter 7
Periodic Properties of the Elements
7.1
Development of the Periodic Table
By the start of the twentieth century, scientists had characterized most of the elements now known to
exist. Various attempts had been made to organize this growing body of information. The best of
these attempts came in 1869 when two scientists, Dmitri Mendeleev and Lothar Meyer, put forth
independent but very similar classification schemes for the elements. Both Mendeleev and Meyer
organized the elements according to physical and chemical properties. Further, they both noticed that
certain properties were repeated periodically if the elements were arranged in order of increasing
atomic weight. Mendeleev went so far as to leave blank spaces in his table where no known element
had the properties that would place it there logically. He called the missing elements eka-aluminum
and eka-silicon, and he predicted their physical and chemical properties based on his theory of
periodic properties. The subsequent discovery of these missing elements and the fact that their
properties so closely match Mendeleev's predictions lent tremendous weight to Mendeleev's theory.
The organization schemes put forth by Mendeleev and Meyer were the forerunners of the modern
periodic table.
In 1913 Henry Moseley improved on the periodic table by developing the concept of atomic
numbers. Arranging the elements in order of increasing atomic number eliminated some of the
inconsistencies in the periodic table, which has been based on atomic weights. For example, in a
table ordered according to atomic mass, potassium would appear below neon in group 8A. Clearly
the properties of potassium, which is a solid at room temperature and highly reactive, would not
logically place it in that group. When the elements are instead arranged in order of increasing atomic
number, argon and potassium appear in their correct places. The modern periodic table arranges
elements in order of increasing atomic number
7.2
Electron Shells and the Sizes of Atoms
According to the quantum mechanical model of the atom, radial electron density—the probability of
finding an electron at a particular distance from the nucleus—falls off as distance from the nucleus
increases, but never really reaches zero. For the noble gases, which exist as isolated atoms, this
complicates the problem of defining atomic "size."
For the other elements in the periodic table, we can determine the radius of an atom by measuring
bond distances between atoms in molecules. For instance, the distance between iodine nuclei in I2 is
2.66 Å. This distance is taken to be the sum of two iodine radii. Therefore, the radius of an iodine
atom is taken to be 1.33 Å. Likewise, the distance between hydrogen nuclei in H2 is 0.74 Å. The
radius of a hydrogen atom is then 0.37 Å. Using these two atomic radii, we would predict a bond
length of 0.37 Å plus 1.33 Å, or 1.7 Å for HI. The measured HI bond length is very close to 1.7 Å.
Using this method, we can determine radii for all of the elements and examine the periodic trends in
atomic size. Two trends emerge.
1. Atomic Radius increases top to bottom within a column
2. Atomic Radius decreases left to right within a period
7.3
Ionization Energy
The minimum amount of energy required to remove an electron from the ground state of an atom or
ion in the gas phase is called the ionization energy. The first ionization energy of magnesium is
represented by the equation:
Magnesium's second and third ionization energies are represented by
and
The first ionization energy (I1) of magnesium is 738 kJ/mol. The second ionization energy (I2) is
1450 kJ/mol. Note that the amount of energy required increased for the second ionization. Because
of coulombic attraction, it is more difficult to remove an electron from a positively charged ion than
it is to remove one from a neutral atom. This third ionization energy of magnesium is 7730 kJ/mol.
This represents a significant increase over the first two. The reason for this jump in ionization energy
is that magnesium has only two valence electrons. The third ionization of magnesium removes an
inner-shell, or "core," electron
So the trends in first ionization energy are
1. First ionization energies generally increase from left to right along the periodic table
2. First ionization energies generally decrease from top to bottom
7.4
Electron Affinities
Essentially the opposite of ionization, electron affinity is the energy required for the addition of an
electron to an atom in the gas phase.
Electron affinity increases as the value of
E becomes more negative
Like the trend across the periodic table for first ionization energy, the trend in electron affinity is not
smooth. Again, the irregularities can be explained in terms of electron configurations.
1. Electron affinities tend to increase (become more negative) from left to right across the
periodic table.
2. Electron affinities tend to remain fairly constant from top to bottom within a group.
7.5
Metals, Nonmetals, and Metalloids
In Chapter 6 you saw how the periodic table could be divided into metals, nonmetals, and
metalloids.
Metals and nonmetals are distinguished by a collection of physical and chemical properties.
Metals tend to lose electrons to become cations. Many transition-metal ions have more than one
possible charge. Figure 7.11 gives the common charges on some transition-metal ions.
Figure 7.11. Charges of some common ions found in ionic compounds.
Notice that the steplike line that divides metals from nonmetals also separates
cation's from anions. Ions shown in yellow have the same number of electrons
as the nearest noble-gas atom.
7.6
Group Trends for the Active Metals
Elements in the same group of the periodic table tend to have very similar physical and chemical
properties. The members of group 1A are called the alkali metals. The alkali metals are soft,
lustrous, very reactive solids.
The chemistry of the alkali metals is very similar. Each readily loses one electron to become a
cation. Each reacts identically with hydrogen, sulfur, and the halogens.
Alkali metals do differ somewhat in their reactions with gaseous oxygen. This is the result of
differences in reactivity. The heavier alkali metals are more reactive, owing to their lower ionization
energies.
The alkaline earth metals (group 2A) exhibit similar physical properties and chemical behavior,
although not as many as the alkali metals do. All alkaline earth metals readily lose two electrons.
The heavier members of the group are more reactive than the lighter ones because of the difference
in ionization energies.
7.7
Group Trends for Selected Nonmetals
Hydrogen is in a group by itself. Its chemistry is dominated by the complete lack of nuclear
shielding of its single electron. It exists as a colorless diatomic gas. Hydrogen reacts with halogens
to form gaseous binary compounds such as HCl. It also reacts with active metals to form solid metal
hydrides such as NaH.
The nonmetallic members of the oxygen group (group 6A) tend to gain two electrons to become
anions. The group contains one metalloid (tellurium) and one metal (polonium). Some properties of
the group 6A elements are given in Table 7.6.
Oxygen plays an important role in terms of the chemistry of other groups of the periodic table.
Oxides of metallic elements, as you have seen, are basic compounds. Oxides of nonmetallic
elements are acids.
Members of group 7A are called the halogens. Halogen chemistry is dominated by the tendency to
gain one electron to become an anion. Halogens in general, and fluorine in particular, are very good
at removing electrons from other species. Because of this, they are often used as industrial oxidants.
Some of the properties of the group 7A elements are given in Table 7.7.
Members of group 8A are known as the noble gases. The noble gases are unique in that they exist as
isolated atoms. They have electron configurations that are associated with exceptional stability and
therefore, as a rule, do not react. Heavier members of the group have been shown to form a limited
number of compounds such as XeF4.
Some properties of the noble gases are given in Table 7.8.