Review for Chapter 8: Periodic Relationships Among the Elements
1. Early versions of the periodic table of elements had the elements arranged according to atomic mass rather than
atomic number. This was not satisfactory in grouping and explaining the characteristics of the elements. For
example, it put argon (atomic mass = 39.95 amu, atomic number = 18) in the place of potassium (atomic mass =
39.10 amu, atomic number = 19) in the periodic table because argon has a greater mass (due to more neutrons) but
lower number of protons than potassium.
2. Arranging the elements by atomic number gave a much more useful periodic table.
3. The properties of an element are determined by its electron configuration. The periodic table is arranged by
atomic number, which also gives the electron configuration.
4. Valence electrons are the outer electrons of an atom and are the ones involved in chemical bonding. The
valence electron configuration directly affects the properties of the atoms of the representative elements.
5. Representative elements are the elements in Groups 1A through 7A. These elements have incompletely filled s
or p subshells of the highest principal quantum number, n, for the period.
6. Noble gases are Group 8A elements. With the exception of helium, they have a completely filled p shell. They
are not reactive because the outer ns and np subshells are completely filled, a condition that represents great
stability.
7. Transition metals are the elements in Groups 1B and 3B through 8B, which have incompletely filled d
subshells or produce cations with incompletely filled d subshells.
8. The Group 2B elements (sometimes labelled 12B) (Zn, Cd, Hg) are neither representative elements nor
transition metals and have no special name. Their d subshell is completely filled.
9. The lanthanide and actinide series of elements are sometimes called f-block transition elements because the
f subshells are being filled.
10. In the formation of cations, one or more electrons are removed from the highest occupied n shell. The
electron configuration for some atoms and their cations are:
Na: [Ne]3s1
Na+: [Ne]
Ca: [Ar]4s2
Ca2+: [Ar]
Al: [Ne]3s23p1 Al3+: [Ne]
Metals in Groups 1A, 2A, and 3A tend to lose electrons until their electron configurations match that of the noble
gas that precedes them in the periodic table as shown above.
11. In the formation of cations from transition metals, electrons are removed from the ns orbital before the
(n-1)d orbital even though the orbitals are filled in the reverse order. For example, Mn, which has an electron
configuration of [Ar]4s23d5, forms a cation, Mn2+, in which two 4s electrons are removed to form [Ar]4s 03d5 instead
of losing two 3d electrons to form [Ar]4s23d3.
12. In the formation of anions, one or more electrons are added to the highest partially filled n shell. The
electron configuration for some atoms and their anions are:
F: 1s22s22p5
F-: 1s22s22p6 or [Ne]
2 2
4
O: 1s 2s 2p
O2-: 1s22s22p6 or [Ne]
2 2
3
N: 1s 2s 2p
N3-: 1s22s22p6 or [Ne]
The nonmetals in Groups 5A, 6A, and 7A tend to accept electrons until their electron configurations match that of
the noble gas that immediately follows them in the periodic table as shown above.
13. Isoelectronic atoms and ions have the same number of electrons and thus the same ground-state electron
configuration. For example F-, Ne, and Na+ are isoelectronic with an electron configuration of 1s22s22p6.
14. The concept of effective nuclear charge can be used to explain variations in atomic size and ionization energy.
For atoms with 3 or more electrons, the electrons in a given shell are shielded from the positive charge of the
protons in the nucleus by electrons in inner shells but not as much by electrons within the same subshell. For
representative elements, the effective nuclear charge increases from left to right across a period and from bottom to
top in a group.
15. The atomic radius is half the distance between two nuclei in adjacent metal atoms or in a diatomic molecule.
Within a period, the atomic radius generally decreases from left to right (Li > Be > B > C > N > O > F > Ne).
Within a group, the atomic radius generally increases from top to bottom (H < Li < Na < K < Rb < Cs). Refer
to Figure 8.5.
16. The ionic radius is the radius of a cation or an anion. If an atom forms a cation, the ion is smaller than the
original atom because one or more electrons have been removed and the electron cloud decreases in size. If an atom
forms an anion, the ion is larger than the original atom because one or more electrons have been added and the
electron-electron repulsion enlarges the electron cloud.
17. For ions, the ionic radius increases from the top to the bottom of the periodic table within a group. For ions of
elements in different groups, a size comparison is meaningful only if the ions are isoelectronic. For example, the
ions Na+, Mg2+, and Al3+ are isoelectronic and decrease in ionic radius from left to right across the period:
Na+ > Mg2+ > Al3+. This is due to an increase in effective nuclear charge moving from Na+ to Al3+. Refer to Figure
8.9.
18. Within a group on the periodic table, physical properties tend to vary predictably, particularly if the
elements exist in the same physical state. For example, the melting point of an element can be estimated from
averaging the melting points of the elements immediately above and below it in its group.
19. The ionization energy is the minimum energy required to remove an electron from a gaseous atom in its ground
state. The ionization energy is a measure of how “tightly” an electron is held in an atom. Ionization is always an
endothermic process.
20. For an atom with more than one electron, each electron has a different ionization energy. The first ionization
energy is the energy required to remove the first electron from the atom in its ground state:
X(g) + energy(I1)  X+(g) + e- . The second ionization energy is the energy required to remove the second
electron: X+(g) + energy(I2)  X2+ + e- . The third ionization energy is the energy required to remove the third
electron: X2+ + energy(I3)  X3+ + e-, and so on.
21. More energy is required to remove each successive electron from an atom since the nuclear charge remains
constant and the effective nuclear charge increases. Ionization energies always increase: I 1 < I2 < I3 < …
22. The alkali metals (Group 1A) have the lowest first ionization energies. The noble gases (Group 8A) have the
highest first ionization energies. First ionization energies tend to decrease from the top of a group to the bottom.
First ionization energies tend to increase across a period from left to right but some exceptions exist. See Figure
8.11.
23. Electron affinity is a measure of the tendency of an atom to gain an electron. The definition states that it is
the negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form
an anion:
F(g) + e-  F-(g) For this reaction, ∆H = -328 kJ/mol, and thus the electron affinity is + 328 kJ/mol.
24. Halogens (Group 7A) have the highest electron affinities. Electron affinity tends to increase from left to right
across a period with some exceptions but varies little within a group. See Figure 8.12.
25. Diagonal relationships are similarities in the chemical properties of pairs of elements that are diagonal to each
other in the periodic table such as Li with Mg, Be with Al, and B with Si. This is due to a similarity in the charge
densities of their cations.
26. The metallic character of elements decreases across a period from metals through the metalloids to the
nonmetals. The metallic character increases from top to bottom within a group of representative elements.
27. Metallic oxides are usually basic, nonmetallic oxides are usually acidic, and oxides of intermediate elements
(such as Al2O3) tend to be amphoteric (displaying both acidic and basic properties).