Periodic Chart
Diderot's Alchemical Chart of Affinities (1778)
Where to begin
• Dobereiner observes
similarities in
elements. Proposes
Law of Triads-Middle
element in the triad
had atomic weight that
was the average of the
other two members.
First Periodic Table
1862
• Alexandre Beguyer de Chancourtois (18201886), professor of geology at the School of
Mines in Paris, publishes a periodic table
constructed as a helical graph.
• Ignored
Helix Periodic Table
John Newlands proposes Octet rule
having arranged the 62 known elements in order of increasing
atomic weights, noted that after interval of eight elements
similar physical/chemical properties reappeared.
• Because these properties seemed to repeat every eight elements,
Newlands called this pattern the law of octaves.
More Periodic History
• Mendeleev
• Then in 1869, Russian
chemist Dimitri
Mendeleev (1834-1907)
proposed arranging
elements by atomic
weights and properties
(Lothar Meyer
independently reached
similar conclusion but
published results after
Mendeleev).
Mendeleev noticed that the chemical
properties of the elements repeated
each time he started a new row.
• Mendeleev made two interesting observations
1. Mendeleev’s table contains gaps that elements with
particular properties should fill.
2. The elements do not always fit neatly in order of atomic mass.
• Mendeleev predicted the properties of the missing
elements.
Mendeleev’s Periodic Chart
More Mendeleev
• Mendeleev's periodic table
of 1869 contained 17
columns with two partial
periods of seven elements
each (Li-F & Na-Cl)
followed by two nearly
complete periods (K-Br &
Rb-I).
Mendeleev Periodic Table
Mendeleev’s Periodic Law
• Properties of the elements is a periodic
function of their atomic mass
Why Mendeleev’s basic version
of the periodic chart?
• Mendeleev/Meyer periodic chart is the basis
of today’s chart because it:
– Predicts Atomic Properties
– Indicates Trends
– Indicates groups that will react/not react
together
Mosley Periodic Table
• After Rutherford’s experiments
• Henry Moseley (1887-1915) subjected
known elements to x-rays. He was able to
derive the relationship between x-ray
frequency and number of protons.
Modern Periodic Law
• Mosley’s Periodic Law- Properties of an element
are a periodic function of their atomic number.
Mosley’s Periodic Table
•When the elements were arranged by increasing
•atomic number, the discrepancies in
•Mendeleev’s table disappeared.
• Moseley’s work led to both the modern
definition of atomic number, and showed
that atomic number, not atomic mass, is the
basis for the organization of the periodic
table.
Last Major Change
• Gene Seaborg
• Starting with plutonium in 1940, Seaborg
discovered transuranium elements 94 to 102 and
reconfigured the periodic table by placing the
lanthanide/actinide series at the bottom of the
table. In 1951 Seaborg was awarded the Nobel
Prize in chemistry and element 106 was later
named seaborgium (Sg) in his honor.
Modern Periodic Law
• Mendeleev’s principle of chemical periodicity
is known as the periodic law, which states:
• The elements are arranged according to their
atomic numbers, elements with similar
properties appear at regular intervals.
A Little More about Mendeleev
and the Periodic Table
S.E. Van Bramer, 9/11
Modified 7/22/99
1
F Block Placed in the Periodic Table
H
3
Li
1
H He
4
5
Be
B C
11 12
6 7
8
9 10
N
O
F Ne
13 14 15 16 17 18
Na Mg
19 20
2
Al Si P
S
Cl Ar
K Ca
21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36
Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
37 38
Rb Sr
39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54
Xe
Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I
55 56 57 58 59 60 61 62 63 64 65 66 67 68 69 70 71 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86
Cs Ba La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
87 88 89 90 91 92 93 94 95 96 97 98 99 100 101 102 103 104 105 106 107 108 109 110 111 112
Fe Ra Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Rf
Db Sg Bh Hs Mt
114
116
118
Circular Periodic Table
Kansas Board Approved Periodic
Table
4 Main Blocks
More Periodic Table
• Period - rows of elements with the same
energy level
• Groups - elements in the same column
with the same electron configuration
• Metals- left of metalloids
• Nonmetals- rt of metalloids
Majority of elements are metals
Special names for
Some Main Groups
• Four groups within the main-group elements have
special names. These groups are:
• alkali metals (Group 1)
• alkaline-earth metals (Group 2)
• halogens (Group 17)
• noble gases (Group 18)
Elements in Group 1 are called alkali
metals.
lithium, sodium, potassium, rubidium,
cesium, and francium
• Alkali metals are so named because they are metals
that react with water to make alkaline solutions.
• Because the alkali metals have a single valence
electron, they are very reactive.
• In losing its one valence electron, potassium
achieves a stable electron configuration.
• Alkali metals are never found in nature as pure
elements but are found as compounds.
Group 2 elements are called
alkaline-earth metals.
• The alkaline-earth metals are slightly less reactive
than the alkali metals.
• They are usually found as compounds.
• The alkaline-earth metals have two valence electrons
and must lose both their valence electrons to get to a
stable electron configuration.
• It takes more energy to lose two electrons than it takes to
lose just the one electron that the alkali metals must give
up to become stable.
Elements in Group 17 of the periodic
table are called the halogens.
• The halogens are the most reactive group of
nonmetal elements.
• When halogens react, they often gain the one
electron needed to have eight valence electrons, a
filled outer energy level.
• Because the alkali metals have one valence electron,
they are ideally suited to react with the halogens.
• The halogens react with most metals to produce
salts.
Group 18 elements are called the
noble gases.
• The noble gas atoms have a full set of electrons in
their outermost energy level.
• The low reactivity of noble gases leads to some
special uses.
• The noble gases were once called inert gases
because they were thought to be completely
unreactive.
• In 1962, chemists were able to get xenon to react,
making the compound XePtF6.
• In 1979, chemists were able to form the first
xenon-carbon bonds.
Hydrogen
• Hydrogen is in a class by itself in the periodic table
• It is placed in group 1 because it has a 1+ charge
and a 1s1 electron configuration
• It can also be placed in Group 17 because of its
behavior as well
Periodicity
• With increasing atomic number, the
electron configuration of the atoms display
a periodic variation.
• Which leads us to Trends
Trends
• Atomic radii
– One-half the diameter of the distance
between two like nuclei
– Increases down a group
– Decreases across a period
Different Atomic Radii
Ionic Radii
Radius of the ion after it has lost or gained an electron
- increases across the chart
- increases down the chart
Ionization Energy
• Amount of energy required to remove an electron
from an atom
– Increases across the chart
– Decrease down the chart
Electron shielding
• is the reduction of the attractive force
between a positively charged nucleus and
its outermost electrons due to the
cancellation of some of the positive charge
by the negative charges of the inner
electrons.
Electron Affinity
• The energy emitted upon the addition of an
electron to an atom or group of atoms in a
gas phase
– Generally become more negative as you go
across the chart\
– No clear trend
Other trends
• Increase in Atomic Number and Mass
• Melting and Boiling Points
-metals increase L-R then decrease to nonmetals
• Electronegtivity
– Energy required by an element in a compound to attract
an electron
- increase from L-R
- decrease top to bottom
Trends in Melting Point