Download CHAPTER 5, THE PERIODIC LAW Section 1, History of the Periodic

CHAPTER 5, THE PERIODIC LAW
Section 1, History of the Periodic Table
By 1860, more than 60 elements had been discovered. FYI, there was no method
for accurately determining an element’s atomic mass or the number of atoms of an
element in a particular chemical compound.
In 1869, Dimitri Mendeleev published his periodic table of elements in which
elements with similar properties were grouped together. The elements were
arranged in order of increasing atomic mass. He left empty spaces in the table. In
1871, he predicted the existence of elements that would fill these empty spaces.
By 1886, all three elements had been discovered: Sc, Ga, and Ge. And, the
properties were strikingly similar to those predicted by Mendeleev. Furthermore,
Mendeleev is considered the discoverer of the periodic law.
In the early 1900s, the chemistry of the lanthanides was understood. The
lanthanides are the 14 elements with atomic number from 58 to 71. To follow was
the discovery of the actinides. The actinides are the 14 elements with atomic
number 90 to 103.
Periodicity with respect to atomic number can be observed in any group of
elements in the periodic table.
Section 2, Electron Configuration and the Periodic Table
Sublevel Blocks of the Periodic Table (image credit, modified:
http://creationwiki.org/Periodic_table_blocks)
KEY
Mendeleev’s Periodic Table (image credit: http://www.chemistryexplained.com/NyPi/Periodic-Table.html)
AM – alkali
metal
AEM – alkaliearth metal
TM – transition
metal
M/S –
metalloid/semimetal
H – halogen
NG – noble gas
L – lanthanide
A – actinide
In 1911, Henry Moseley rearranged the periodic table of elements by putting
elements in order of increasing nuclear charge, or number of protons in the
nucleus. His work led to the modern definition of atomic number and the
recognition that atomic number, not atomic mass, is the basis for the organization
of the periodic table.
Periodic law states that the physical and chemical properties of the elements are
periodic function of their atomic numbers.
The periodic table is an arrangement of the elements in order of their atomic
numbers so that elements with similar properties fall in the same column, or group.
Between 1895 and 1898, Sir William Ramsay proposed adding the Group 18
elements, the noble gases, to the periodic table.
Periodic Table of Elements (image credit, modified:
http://intro.chem.okstate.edu/AP/2003Norman/HTML/SCH7.HTM)
s-block
• Group 1, alkali metals
• Group 2, alkaline-earth metals
• Hydrogen
• Helium (noble gas)
p-block
• Groups 13-18
o Metalloids or Semiconductors
o Group 17, halogens
o Group 18, noble gases
PERIODIC TRENDS
Moving from left to right
across a period
Atomic
Radii
Decreases
Why? Increasing positive charge of
the nucleus.
Ionization
Energy
Increases
Why? Increasing nuclear charge.
Electron
Affinity
Increases (becomes more negative)
Why? Increasing nuclear charge
allow for electrons to be more easily
acquired.
Ionic
radii
Decreases
Why? The electron cloud shrinks due
to increasing nuclear charge.
Increases
d-block
• Groups 3-12, transition metals
f-block
• Lanthanides
• Actinides
Section 3, Electron Configuration and Periodic Properties
Atomic radius is defined as ! the distance between the nuclei of identical atoms that
are bonded together.
An ion is an atom or group of bonded atoms that has a positive or negative charge.
Ionization is any process that results in the formation of an ion.
The atom’s electron affinity is the energy change that occurs when an electron is
acquired by a neutral atom.
A cation is a positive ion. The formation by loss of one or more electrons ALWAYS
leads to a decrease in atomic radius because the removal of the highest-energy-level
electrons results in a smaller electron cloud. And, the remaining electrons are drawn
closer to the nucleus by its unbalanced positive charge.
An anion is a negative ion. The formation by the addition of one or more electrons
ALWAYS leads to an increase in atomic radius. This is because the total positive
charge of the nucleus remains unchanged when an electron is added to an atom or an
ion. So, the electrons are not drawn to the nucleus as strongly as they were before the
addition of the extra electrons. The electron cloud also spreads out because of greater
repulsion between the increased number of electrons.
Electronegativity is a measure of the ability of an atom in a chemical compound to
attract electrons from another atom in the compound.
Electronegativity
Moving from top to bottom
down a group
Increases
Why? Electrons occupy sublevels in
successively higher main energy
levels located farther away from the
nucleus.
Decreases
Why? Additional energy levels and
electron shielding. Electrons are
farther from nucleus. Outer
electrons are partially shielded from
the effect of the nuclear charge by
inner electrons.
Decreases (becomes less negative)
Why? Increasing atomic radius
outweighs increasing nuclear
charge, causing addition of electrons
difficult.
Increases
Why? Occupationof successively
higher energy levels.
Decrease or remain about the same