Periodic Table
Chapter
4.1, 4.2, 4.3
4-1: Development of the Periodic Table
1) History of the Periodic Table – By the end
of the 1700’s scientists had identified only 30
elements
(ex. Cu, Ag, Au, H2, N2, O2, C). By the mid
1800’s by using spectroscopy additional
elements were identified by using their line
spectra and about 65 elements had been
identified.
A. J.W. Dobereiner: 1829
Organized the elements into
groups with similar properties.
He called these groups triads.
The middle element is often
the average of the other two.
Ex) Cl – 35.5
Cl + I
 Avg.
Br – 79.9
2
I – 126.9
Ca
Avg
Sr
Ba
Triads on the Periodic Table
B. J.A.R. Newlands: 1867
• Law of Octaves.
He said properties
repeated every 8th
element. There
were 62 known
elements at the
time. He was also
a musician.
C. Dimitri Mendeleev: 1869
1. Organized the 1st
periodic table
according to
increasing atomic
mass and put
elements with
similar properties in
the same column.
2. He arranged some
elements out of atomic
mass order to keep them
together with other
elements with similar
properties. He also left
three blanks in his table
and correctly predicted
the properties of these 3
unidentified elements
that were later
identified and matched
his predictions.
D. Moseley: 1915
Each element has a certain amount
of positive charge in the nucleus
which are called protons.
1. Moseley reorganized the periodic
table by Atomic Number.
2. The Periodic Law: When elements
are arranged in order of increasing
atomic number, their physical and
chemical properties show a periodic
pattern.
Glenn Seaborg “Seaborgium”
Sg #106
• Born in 1912 in Michigan,
Seaborg proposed
reorganizing the Periodic
Table one last time as a
young chemist working on
the Manhattan Atomic Bomb
Project during WWII by
pulling the “f-block” elements
out to the bottom of the table.
He was the principle or codiscoverer of 10 transuranium
elements. He was awarded
the Noble prize in 1951 and
died in 1999.
4-2: Reading the Periodic Table
A. Information in each square:
Atomic # (protons)
Mass # - protons & neutrons
particular isotope of
aluminum like
Al-26 or Al-27
Symbol
13
Al
Name
Aluminum
Atomic Mass
26.9815
Weighted average of all an element’s isotopes
B. Organizing the Squares
1. Vertical Columns – groups & families
2. Horizontal Rows - periods
Parts of the Periodic Table
C. Electron Configuration & Families
1. Valence electrons –
outermost electrons
responsible for
bonding.
2. Elements in the
same group have
the same number
of valence
electrons.
Carbon has 4 valence electrons
E. Metal, Nonmetals and Metalloids (Semimetals):
1. Metals
- good conductors of heat & electricity
- most solids at room temperature (except Hg)
- high luster (shiny)
- ductile (can be drawn into thin wire)
- malleable (bends without breaking)
- high melting points
- high densities
- react with acids
2. Nonmetals
- brittle (shatters when struck)
- low luster (dull)
- neither ductile nor malleable
- nonreactive with acids
- nonconductors
3. Metalloids - Properties of both metals & nonmetals
(Semimetals)
Alkali Metals
have 1 valence
electron:
Brainiac Alkali
Metal Video
English Video –
Noble Gases Clip
4-3: Trends in the Periodic Table
A. Atomic Radius
1. The distance from
the center of the
nucleus to the
outermost electron.
2. Atoms get larger
going down a group
and smaller going
across a period.
Ex) Na is larger than Mg
Na is smaller than K
Atomic Radii of the
Representative Elements
Atomic Radii vs Atomic Number
Ionic Size
1. When atoms gain
electrons, they
become (-) and get
larger.
Positive Ion Size
1. When atoms
lose electrons,
they become
(+) and get
smaller.
2. Ions get larger
as you go
down a group.
Relative Sizes of Positive &
Negative Ions
The sodium ion lost an electron, and
therefore the positive protons in the
nucleus exert a stronger pull on the
remaining negative electrons,
shrinking the orbitals. Thus positive
ions are smaller than their atoms.
The chloride ion gained an electron, and
therefore the fewer positive protons in
the nucleus exert a weaker pull on the
extra negative electrons, increasing the
size of the orbitals. Thus negative ions
are larger than their atoms.
Electron Attraction in a Bond &
Ion Size
C. Ionization Energy:
1. The energy needed to remove one electron
from an atom.
2. Elements that do not want to lose their
electrons have high ionization energies.
3. Elements that easily lose electrons have low
ionization energies.
4. I.E. decreases down a group (opposite of
atomic radius)
5. I.E. increases across a period. (opposite of
atomic radius)
Ex) Na IE smaller than Mg
Na IE larger than K
Ionization Energy of the
1st 20 Elements
Ionization Energy vs. Atomic Number
Sublevels by the Periodic Table
D. Successive Ionization Energies:
1.
2.
Energy required to remove electrons beyond the 1st electron.
Ionization energies will increase for every electron removed.
4560 kJ
3. Na [Ne]3s1 Na• 1st = ____
496 kJ 2nd = ____
738 kJ 2nd = ____
7730
1450 kJ 3rd = ____kJ
4. Mg [Ne]3s2 Mg: 1st = ____
11,600
1816 3rd = 2744
577 2nd = ____kJ
5. Al [Ne]3s23p1 Al: 1st = ____kJ
____kJ 4th = ___kJ
E. Electron Affinity:
1. Energy change that occurs when an atom gains
an electron.
2. A highly negative electron affinity attracts
electrons. (nonmetals)
3. A positive electron affinity does not attract
electrons. (metals)
Electron Affinity
decreases
F. Electronegativity:
1. Reflects an atoms ability to attract electrons
in a chemical bond.
2. E.N. decreases going down a group
3. E.N. increases going across a period.
4. Examples: NaCl and H2
G. Octet Rule:
1. An atom will tend to lose, gain or share
electrons in order to acquire a full set of
valence electrons.
2. “Octet” = 8 = s2p6 configuration
H. Oxidation Number:
The charge on an ion when it gains or
loses electrons to acquire a stable octet.
Which of the following would have the largest?
• Atomic Radii?
• Ionization Energy?
• Electron Affinity?
Element D
Element C
Element A
Which of the following would have the smallest?
• Atomic Radii?
• Ionization Energy?
• Electron Affinity?
Element B
Element D
Element D
Electron Shielding