Chapter 5: Periodic Law
I.
Worksheet on History of the Periodic Table and Periodic Law.
A. Dobereiner, Newlands, Mendeleev, Moseley.
B. Periodic Law: The physical and chemical properties of the elements
are periodic functions of their atomic numbers.
II. The Periodic Table
A. Periodic Table: an arrangement of the elements in order of their
atomic numbers so that elements with similar properties fall in the same
column or group.
B. Information in each block of the table:
1. Symbol
2. Atomic Number
3. Atomic Mass
4. Name
5. Electron Configuration
C. The Periodic Table is organized into rows and columns of elements.
1. Periods: Horizontal rows of elements. Numbered 1-7. Those
numbers correspond to the principle energy level (n).
2. Groups: Vertical columns of elements. Also called families.
numbered in 3 ways.
a. European System: IA-VIIIA; IB-VIIIB
b. U.S. System: 1A, 2A, 3B-8B, 1B, 2B, 3A-8A. We will use this
one.
c. IUPAC System: 1-18
3. Family Names to know:
Group
Family Name
1A
Alkali Metals
2A
Alkaline Earth Metals
7A
Halogens
8A
Noble Gases
Other groups with A are called by the name of the first element in
that group. For example, Group 4A is called the Carbon Group
(Family).
4. The elements with B in the group number are the transition metals.
5. The Lanthanides and Actinides are on the bottom of the table. The
are called the inner transition metals. They don’t fit into the table
very well. They make it too big.
6.
Elements are also classified into Metals, Nonmetals, and Metalloids
(Semimetals.) Dividing line starts at B and zigzags down the table.
III. Electron Configuration and the Periodic Table
A. Blocks of the Periodic Table
1. Notice that all elements in the same group have configs that end
similarly. For example group 1A all end in Xs1. Group 4A all end in
Xs2 Xp2. We use this information to break the table down into
blocks.
2. Groups 1A and 2A are the s block. 3A-8A are the p block. TM’s are
the d block. ITM’s are the f block. If given an element, you should
be able to tell me the block that it is in.
3. We can use this information, group numbers, and period numbers
to help us get our configs much faster.
4. Practice with Ag, Pd, Nb, Ca, S, and I. Note that exceptions to the
Aufbau Principle still are exceptions.
IV. Periodic Properties
A. Atomic Radii
1. Atomic Radius is defined as one-half of the distance between the
nuclei of identical atoms that are bonded together.
2. AR is measured (in pm) this way because the boundary is fuzzy and
varies under different conditions.
3. As you move down a group of elements AR will increase. This
occurs because for every period that you move down, you are
adding another principle energy level(ring of Bohr Model).
4. As you move across a period of elements AR will decrease. This is
because of the stronger attraction between the oppositely charged
particles in the atom. Compare two weak magnets with two strong
magnets. Figure 13 on page 151.
B. Ionization Energy
1. Ionization Energy is the energy required to remove 1 electron
from a neutral atom of an element.
2. IE is measured (in kJ/mol) on individual atoms in the gas phase, to
eliminate interference.
3. It is a measure of how strongly an atom holds onto its outermost e-.
4. IE decreases as you move down a group. Further away from the
nucleus, held more loosely.
5. IE increases as you move across a period. Closer to nucleus, held
tighter. Figure 15 on page 153.
6.
It is possible to remove electrons from a positive ion. It will require
more energy. This would be the second (third, fourth, etc)
ionization energy. More energy would be needed because the
electrons get held in tighter after the first ionization. Refer to Table
3 on page 155.
C. Electron Affinity
1. Electron Affinity is the energy change that occurs when an
electron is acquired by a neutral atom.
2. Because most atoms release energy when this happens, EA is
measured in - X kJ/mol.
3. Trends for EA do not change as regularly as IE trends. Halogens
tend to have the most negative EA. Refer to Figure 17 on page 157.
D. Ionic Radii
1. Ionic Radius is similar to AR in terms of what is being measured.
2. When an atom gains an electron it becomes negatively charged.
Negatively charged ions are called anions.
3. When an atom loses an electron it becomes positively charged.
Positively charged ions are called cations.
4. Anions will be larger than their respective neutral atoms. More
electrons need more room to move around. There is also repulsion
between the electrons.
5. Cations will be smaller than their respective neutral atoms. When
the outer electron is removed, the nucleus can pull harder on the
others, making the ion smaller.
6. See Figure 19 on page 159.
E.
Valence Electrons
1. Valence electrons are the outermost electrons. When bonding
occurs, it is involving valence electrons.
2. For our purposes, all elements in a group with an A designation will
have the number of valence electrons as the group number. For
example, potassium, K, is in Group 1A. It has 1 valence electron.
F.
Electronegativity
1. Electronegativity is a measure of the ability of an atom in a
chemical compound to attract electrons.
2. Fluorine is the most EN element. It is assigned an arbitrary value of
4.0. All other EN’s are calculated with respect to F.
3. EN tends to increase as you move across a period. Nonmetals want
electrons. Therefore they have higher EN values.
4. EN tends to decrease as you move down a group.
5.
6.
Noble gases are not assigned EN values because they do not
normally form bonds.
High EN elements are generally in the upper right of table. Low EN
elements are generally in the lower left of table. Figure 20 on page
161.