SCIENCE AND
TECHNOLOGY III
CHEMISTRY
Learning
Content:
·
Trends
in the Periodic Table
§ Sizes of
Atoms and ions
§ Ionization
Energy
§ Electron
Affinity
§ Electron
Negativiy
References:
[1] SETS- Chemistry by Ungson
[2] General chemistry by Petrucci
Lesson Proper: Part I-Sizes of Atoms and ions
PERIODIC
TRENDS
1. Atomic Radius (Atomic
Size)
- Unfortunately, atomic radius is hard to define. However, we can measure
the distance between the nuclei of adjacent atoms (internuclear distance).
Thus, we can emphasize an atomic radius based on the distance between the
nuclei of two atoms joined by a chemical bond.
- Types of atomic radii include covalent
radii, metallic radii, and ionic radii.
§
Covalent radius- is one-half the distance
between the nuclei of the two atoms joined by a single covalent bond.
§ Metallic radius- as one-half the distance between the nuclei of two
atoms in contact in the crystalline solid metal.
§ Ionic radius- is based on the distance between the nuclei
of ions joined by an ionic bond. Because the ions are not identical in size,
this distance must be properly apportioned between the cation and anion. One
way to apportion the electron density between the ions is to define the radius
of one ion and then infer the radius of the other ion.
- Variation of Atomic Radii Within a Group
of the Periodic Table.
Moving
down a group, the principal quantum number (n) of the outermost electrons
increases. Electrons with a larger principal quantum number are found in
orbitals that extend farther away from the nucleus, which makes the atomic
radius larger.
Nevertheless, the more electronic shells
in an atom, the larger is the atom.
Atomicradius increases from top to bottom through a group
of elements.
- Variation
of Atomic Radii Within a Period of the Periodic Table.
In
any period, the outer electrons of each element are in orbitals with the same
principal quantum number. But moving from left to right across a period, the
atom’s nuclei gain more protons. Atoms that have a more positive charge in their
nuclei exert a stronger pull on the electrons in a given principal quantum
number. Thus, the atom’s outer electrons are increasingly attracted to the
nucleus. It shrinks the electrons’ orbitals and makes the atom smaller. The
distance between the nucleus and the electrons does not increase as there is no
added orbital, thus resulting in a stronger attractive force and “shrinking the
atom”.
Thus, the atomic radius
decreases from left to right through a period of elements.
EXAMPLE
1: Relating Atomic Size to Position in the Periodic Table
Refer only to the periodic table on the
inside front cover of your textbook, and determine which the largest atom is:
Sc, Ba, or Se.
Analyze
We first place the element in the periodic
table and decide whether or not the elements are in the same period and whether
they are on the right or left of the periodic table. We can then use the rules
noted above to decide on the relative sizes of atoms (or ions).
Solve
Sc and Se are both in the fourth period, and
we would expect Sc to be larger than Se because atomic sizes decrease from left
to right in a period. Ba is in the sixth period and so has more electronic
shells than either Sc or Se. Furthermore, it lies even closer to the left side
of the table (group 2) than does Sc (group 3). We can say with confidence that
the Ba atom should be the largest of the three.
PRACTICE
EXERCISE 1: Use the periodic table on
the inside front cover of your textbook to predict which is the smallest atom:
As, I, or S.
_______________________________________________________________________________
2. Ionic Size
- When
an atom loses electrons – or becomes a positive ion – it becomes smaller.
Loss
of electrons vacates the atom’s largest orbitals and reduces the repulsive
force
between the remaining electrons, allowing them to be pulled closer
to
the nucleus.It becomes smaller because it has fewer electrons.
The given figure compares four species:
the atoms Na and Mg and the ions Na+
And Mg2+. As expected, the Mg atom is
smaller than the Na atom, and the cations
are smaller than the corresponding atoms.
Na+ and Mg2+ are isoelectronic- they
have
equal numbers of electrons (10) in identical
configurations, 1s22s22p6.
Mg2+ is smaller
than Na+ because
its nuclear charge is larger
(+12, compared with +11 for Na)
·
For isoelectronic cations, the more positive the
ionic
charge, the smaller the ionic
radius.
- When
an atom gains electrons – or becomes a negative ion – it becomes larger.
Greater
number of electrons increases the electric repulsion forces among them causing
the electrons to spread out, thus, making the ion larger.
It
becomes larger because of the increase in the number of electrons.
EXAMPLE 2: Comparing the Sizes of Cations and Anions
Refer only to the periodic
table on the inside front cover of your text book, and arrange the following
species in order of increasing size: K+, Cl-, S2-
and Ca2+.
Analyze
The key lies in recognizing
that the four species are isoelectronic, having the electron
configuration of argon: 1s22s22p63s23p6.
When considering isoelectronic cations, the higher the charge on the ion, the
smaller the ion.
Solve
The larger charge on the calcium
ion means that Ca2+ is smaller than K+. Because K+ has
a higher nuclear charge than Cl- (Z=19 compared Z=17 with), it is
smaller than Cl- . For isoelectronic anions, the higher the charge,
the larger the ion. S2- is larger than Cl- . The order of
increasing size is
Ca2+< K+<
Cl- < S2-
Assess
We can summarize the
generalizations about isoelectronic atoms and ions into a single statement:
Among isoelectronic species, the greater the atomic number, the smaller the
size.
PRACTICE EXERCISE 2: Refer only to
the periodic table on the inside front cover, and arrange the following species
in order of increasing size: Ti2+, V3+, Ca2+,
Br- and Sr2+.
PRACTICE EXERCISE 3: On the blank periodic table below, locate
the following:
(a) The smallest group 13
atom (b)
The smallest period 3 atom
(c) The largest anion of a
nonmetal in period 3 (d)
The largest group 13 cation
