The Periodic Law
History of the Periodic Table
Problems
• Different chemists used different
masses for elements.
• There were no accurate ways to
determine the number of atoms in a
compound.
• Due to different masses, there were
different compositions calculated.
Mendeleev and Chemical Periodicity
• Used values of atomic masses
established by Cannizzaro.
• Wished to organize elements by their
properties.
• Looked for trends or patterns in
properties.
• Found repetitive properties--periodicity---when elements were
arranged according to mass.
Mendeleev and Chemical Periodicity
• Grouped elements with similar
properties together.
• Created the first periodic table.
• Left open spaces for predicted
elements.
• Note: The elements predicted by
Mendeleev in 1871 were found by 1886
and their properties were almost exactly
those predicted by Mendeleev years
earlier.
Properties of Some Elements
Predicted by Mendeleev
Moseley and the Periodic Law
• Organized elements by increasing
nuclear charge.
• Led to the atomic number concept.
• Reorganized the periodic table in order
of atomic number.
• Supported Mendeleev’s periodic law
which stated “The physical and
chemical properties of the elements are
periodic functions of their atomic
numbers.”
The Modern Periodic Table
• It 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.
• The most significant additions were the
noble gases. Helium was discovered in
1868. Argon was discovered by
Ramsey in 1894. He also discovered
krypton and xenon in 1898. Dorn
discovered radon in 1900.
The Modern Periodic Table
• There were additions to the table after
Mendeleev’s time.
• The noble gases---They are known by
this name because they are nonreactive
or inert.
• The lanthanides---elements 58 to 71.
• The actinides---elements 90 to 103.
The Periodic Law
Electron Configurations and the
Periodic Table
Periods and Blocks of the Periodic
Table
• In the first period, the 1s sublevel is
filled. Since it holds only 2 electrons,
the first period must fill with 2
electrons.
• In the second period, there are ___
sublevels, ___ and ___.
• In the third main level, there are ___
sublevels.
Periods and Blocks of the Periodic
Table
• The period of an element can be found
by its configuration. Arsenic, for
example, is [Ar]4s23d104p3 which means
that it is in the fourth period because
the highest occupied main level is n =
4.
• Based on the electron configurations of
the elements, the periodic table can be
divided into four blocks, the s, p, d, and
f blocks.
Blocks of the Periodic Table
The s-Block Elements: Groups 1 and
2
• These are chemically reactive metals. The
Group 1 (ns1) metals are more reactive than
the Group 2 (ns2) metals.
• Group 1 elements contain ___ valence
electron(s) and are known as the alkali
metals.
• Group 2 elements contain ___ valence
electron(s) and are known as the alkaline
earth metals.
• Hydrogen and Helium are in the s-block but
their properties are not like the others in the
block. They are unique.
The d-Block Elements: Groups 3-12
The Transition Metals
• The d-sublevel first appears in the ___ main
level.
• The d-sublevels are slightly higher in energy
than the previous s-sublevel. This is where
the main energy levels begin to overlap.
• There are exceptions in this block. Some
elements fill electrons to achieve the lowest
possible energy state. Two exceptions in the
4th period are chromium and copper.
Experimentally it was found that sublevels
prefer to be half-filled or completely filled if
they are close to either state.
The d-Block Elements: Groups 3-12
• Chromium has the configuration
1s22s22p63s23p64s13d5. Molybdenum would be
similar.
• Copper has the configuration
1s22s22p63s23p64s13d10. Silver would be similar.
• Palladium, platinum, and many others are also
exceptions to the normal filling pattern. In each
case the abnormal filling occurs to provide the
lowest possible energy state for the atom.
• These elements are not as predictable as the sblock or the p-block. They usually have more
than one oxidation state.
Exceptional Electron Configurations
• Niobium [Kr] 5s14d4
• Molybdenum [Kr] 5s14d5
• Technetium [Kr] 5s14d6
• Ruthenium [Kr] 5s14d7
• Rhodium [Kr] 5s14d8
• Palladium [Kr] 4d10
• Silver [Kr] 5s14d10
• Platinum [Xe] 6s1 4f14 5d9
The p-Block Elements: Groups 13-18
• These are atoms which add an electron to the
p-sublevel.
• These are main-group elements or
representative elements.
• All of the nonmetals and metalloids are
included in the p-block. There are 8 metals
here also.
• The halogens(Gp 17) are the most reactive
nonmetals. Fluorine is the most reactive
nonmetal on the periodic table.
The f-Block Elements: Lanthanides
and Actinides
• These are located on the bottom of the
periodic table but are really placed in
the sixth and seventh periods.
• The first main level that has an fsublevel is the ___ level.
• There are 14 elements on each row.
• They are similar in reactivity to Group 2
elements.
Main Group (Representative) Elements
• Group 1 – Alkali metals – very reactive metals
– never found freely in nature
• Group 2 – Alkaline earth metals – reactive
metals – never found freely in nature
• Group 3(13) – Boron group – less reactive –
contains a metalloid and metals
• Group 4(14) – Carbon group – most diverse
group
• Group 5(15) – Nitrogen group
• Group 6(16) – Chalcogens
• Group 7(17) – Halogens – most reactive
nonmetals
• Group 8(18) – Noble gases – most stable
Relationship between Periodicity and
Electron Configuration
Electron Configuration and
Periodic Properties
Trends Exhibited on the Periodic Table
Atomic and Ionic Radii
• The atomic radii of elements decrease
as one moves across a period. (Why?)
• This is caused by an increasing nuclear
positive charge while electrons are
added into the same main level.
• The atomic radii of elements increase as
one moves down a group.(Why?)
• This occurs because more energy levels
are being filled as one goes down a
group.
Trends in Atomic Radii
Atomic Radii
• Going down groups, the radii of transition
elements increase from the 4th to the 5th
period. However, the 6th period transition
elements, are about the same size as the
5th period transition elements.
• This is due to the lanthanide contraction,
the placement of the lanthanides after the
6s sublevel. The expected increase in size
does not occur.
Trends in Atomic Radii
Ionic Radii
• Atoms form ions by losing or gaining
electrons.
• An atom that has lost an electron
becomes ___ and an atom that has
gained an electron becomes ___.
• A positive ion is called a ___.
• A negative ion is called an ___.
Isoelectronic Character
• Isoelectronic means that species will have the
same number and distribution of electrons.
Ions will have the same electron distribution as
the nearest noble gas.
• There is an isoelectronic series on every row.
• The more positive ion in such a series is the
smallest. The more negative ion in such a
series is the largest.
Isoelectronic Series
• An isoelectronic series is Al+3, Mg+2,Na+1, Ne, F-1,
O-2, and N-3. All are 1s22s22p6. Which of the
above ions is largest?
• N-3 is. Why is that so?
• Nitrogen has 7 protons and 7 electrons, but
when the ion is formed, 3 more electrons are
added. There is an imbalance in the positive and
negative charges. There are more negatives
than positives so the electron cloud is not held as
tightly. In this species there are the fewest
protons for the 10 electrons, so it is the largest.
Ionic Radii
Trends in Ionization Energies
Ionization Energy
• Ionization energy is the amount of energy
required to remove an electron from a
gaseous atom. This is always an
endothermic change.
– A + first ionization energy  A+ + e-
• It is usually expressed in kJ/mol.
• The value of the first ionization energy
depends on the effective nuclear charge, the
atomic radius, and the electron
configuration.
Ionization Energy
• I.E. increases with each successive electron
removal. (Why?)
• As each electron is removed there is more
nuclear charge on each electron remaining
with the atom. Therefore more energy is
required.
• The amount jumps dramatically when trying
to remove an electron from a filled, stable
level.
Ionization Energy
• Which of the following species would
require the most energy to remove a
first electron? Rb
or
I
• Iodine because there are more
electrons between the outer electron
and the nucleus but the single electron
of Rb does not have that
protection(inner electron shielding).
• I.E. increases across to the Noble
Gases and decreases down a group.
Electron Affinity
• Neutral atoms can also acquire electrons.
• The energy change that occurs when an
electron is acquired by a neutral atom is
called the electron affinity.
• A + e-  A- + energy or A + e- +
energy  A- may be the general
equation. The change may be
endothermic or exothermic, but is
exothermic most of the time.
Electron Affinity
• Elements in the upper right corner (Gp 17)
tend to have large negative electron
affinities meaning that they become more
stable by gaining an electron. That would
be a high electron affinity.
• Elements in the lower left corner tend to
have much less negative electron
affinities. They are unlikely to gain an
electron.
Electron Affinity
Electronegativity
Electronegativity
• Electronegativity is a measure of an
atom’s ability to attract electrons in a
compound.
• Fluorine is the most electronegative
element on the periodic table.
• Electronegativity decreases as one
moves down and to the left.
Summary of Trends