Download Unit 10 Lecture Notes

Chemistry--Unit 10: Chemical Periodicity
Lecture Notes
I.
II.
Classification of the Elements
A. The Periodic Table Revisited
1. The periodic table is probably the most important tool in chemistry
2. The name “periodic” is associated with the periodic law which states that
the properties of elements are a periodic function of their atomic number
(simply, properties of elements repeat according to the periodic table
arrangement)
B. Classifying Elements by Electron Configuration
1. Both the position of an element on the periodic table and the properties of
that element arise from the electron configuration of the atoms
a. The noble gases (group 18) all have an ending electron configuration
of s2p6, they are all inert, unreactive gases
1) the most stability an atom can have is having a full outer energy
level (full s and p)
2) some stability is also associated with full and half-full sublevels
b. Recall the s, p, d, and f blocks; the last electrons added to the atom put
elements in these blocks
c. The transition metals are transition metals because the last electrons
added are in the d sublevel, the inner transition metals are inner
transition metals because the last electrons added are in the f sublevel
d. The representative elements are those in groups 1 (alkali metals), 2
(alkaline earth metals), 13 (boron family), 14 (carbon family), 15
(nitrogen family), 16 (oxygen family), and 17 (halogens); notice the
group # indicates how many electrons are in the outermost energy
level
2. Since electron configuration follows a pattern, many properties of
elements also follow a pattern based on configuration
a. Elements in the same group have similar chemical and physical
properties
b. Elements in different groups have different properties, even if they are
in the same period
3. Density and boiling point, for example, are "periodic properties"; there are
many others (next part of chapter!)
Periodic Trends
A. Trends in Atomic Size
1. As the principle quantum number increases, the size of the electron cloud
increases
2. Because of this, the size of atoms in each group increases as you move
down the table
3. Atomic radii generally decrease as you move from left to right through
each period since more protons and electrons are being added, the total
attraction between them increases making the atom smaller
4. In summary, atomic radii increase top to bottom and right to left in the
periodic table
Chemistry--Unit 10: Chemical Periodicity
Lecture Notes
B.
C.
D.
E.
5. Noble gases do not follow this trend because in their full octet there is no
electron interaction
6. The size of atoms is determined using X-ray diffraction in which an X-ray
is passed through the crystals of an element
Trends in Ionic Size
1. When positive ions form, the resulting ion is smaller than the neutral
parent atom
a. usually loses an outer energy level
b. more protons attracting fewer electrons
2. When negative ions form, the resulting ion is larger than the neutral parent
atom because more electrons causes more repulsive force within the
electrons and they are pushed further away from each other
3. In general, metallic ions on the left and in the center of the table are
formed by the loss of electrons. They are smaller than the atoms from
which they are formed. Nonmetallic ions are located on the right side of
the table. They are formed by the gain of electrons and are larger than the
atoms from which they are formed
Predicting Oxidation Numbers
1. Recall trend of +1, +2, transition metals of various positive oxidation
states (charges of monatomic ions), +3, 4, ─3, ─2, ─1
2. Also recall stability of the noble-gas configuration, as well as the stability
of full and half-full sublevels
3. Also be able to predict and explain oxidation numbers
Trends in Ionization Energy
1. The energy required to remove an electron from an atom is called its
ionization energy
2. The energy needed to remove the most loosely held electron from an atom
is called the first ionization energy
3. Both are measured in kilojoules/mole
4. Ionization energy is a periodic property
a. The ionization energy tends to increase as atomic number increases in
any horizontal row or period. In any group or column, there is a
gradual decrease in first ionization energy as atomic number increases
b. Both increased distance of the outer electrons from the nucleus and the
shielding effect (in which inner electrons block the attraction of the
nucleus for outer electrons) tend to lower ionization energy as you
move down a group
c. In moving across a period of the periodic table, there is a general
increase in first ionization energy as a result of increasing nuclear
charge
d. Because of the stability of full and half-full sublevels, there are a few
exceptions to the trend in ionization energy across a period
e. A metal is characterized by a low first ionization energy
f. An element with a high first ionization energy is a nonmetal
Multiple Ionization Energies
Chemistry--Unit 10: Chemical Periodicity
Lecture Notes
1. The energy required to remove the second and third (and more) electron
from an atom can be measured resulting in measurements of multiple
ionization energies
2. The second ionization energy would be the energy required to remove the
second most loosely held electron from an atom, the third ionization
energy would be the energy required to remove the third most loosely held
electron, and so on
3. For each subsequent electron removed from an atom, the ionization energy
required increases
4. Trends do exist in multiple ionization energies, and the electron
configuration of the atom must be considered when explaining values
obtained for multiple ionization energies
5. By knowing multiple ionization energies alone for a particular atom,
oxidation states can be predicted; a large jump in the values would
indicate a likely oxidation state
F. Trends in Electronegativity
1. The electronegativity of an element is the tendency for the atoms of the
element to attract electrons when they are chemically combined with
atoms of another element
2. Electronegativity is measured on the Pauling electronegativity scale, it has
units of “Paulings” and ranges from 0.0 to 4.0.
3. The same factors that affect ionization energy also affect electronegativity
4. Electronegativity generally decreases as you move down a group; since
the nucleus is farther from the outside of the atom, the less attraction it has
for bonded electrons, greater shielding effect as well
5. Electronegativity generally increase as you move across a period; the
smaller atom creates a smaller distance between the nucleus and the
bonding electrons, and so a greater attraction between the two (nuclear
charge is also increasing