Periodic Law
Periodic Table
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Prior to 1860 no agreement/method to
accurately determine masses of atoms.
First International Congress of Chemists – 1860
Stanislao Cannizzaro presented method for
accurately measuring atomic masses
 Looked for relationships between atomic masses and
other properties of elements
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First tables arranged elements by atomic weight
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Could not agree on atomic weights therefore tables
were different
John Newlands
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Noticed elements properties repeated every 8th
element when arranged by atomic mass
Named this phenomenon “the Law of Octaves”
Did not work for all elements
Julius Lothar Meyer
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Developed first modern table
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Consisted of 28 elements divided
into 6 families
Families (groups) had similar
chemical and physical properties
Discovered all elements in same
family had same number of
valence e- -- outermost electrons
in highest energy level
Why?
Dmitri Mendeleev
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Noticed that properties repeat
themselves at certain intervals
Arranged all known elements
into one table based on
properties– 1869
1871 - Proposed the “Periodic
Law”
Based on the properties spaces
were left for unknown
elements (Sc, Ga, Ge)
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Upon discovery of other elements
inconsistencies were found with Mendeleev’s
table
Atomic masses improved and they no longer
arranged the elements by increasing atomic mass
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Why can most elements be arranged by atomic
mass?
What was the reason for chemical periodicity?
Henry Mosely
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Discovered elements contain unique number of
protons (atomic number) - 1911
Arranged elements by atomic number - 1913
Fully explained the Periodic Law
Periodic Law
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The physical and chemical properties of the
elements are periodic functions of their atomic
numbers.
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Aka – when elements are arranged by increasing
atomic number, elements with similar properties
appear at regular intervals.
Parts of the Periodic Table
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Noble Gases – added to the table in 1894 after
the discovery by Lord Rayleigh and William
Ramsey
First discovered Argon while studying nitrogen
 Later discovered Helium
 Highly inert (unreactive) due to a full octet
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Parts….
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Alkali metals – group 1
Alkaline earth metals – group 2
Halogens – group 17
Transition metals – d block elements
Inner Transition metals
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Lanthanides (elements 58-71) added in early 1900’s
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Have very similar properties
Actinides (elements 90-103)
Electron Configuration
& the Periodic Table
s-Block Elements
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Groups 1 & 2
All elements in group 1 & 2 will have an
electron configuration of
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ns1 or ns2 where n = highest energy level occupied
Alkali Metals
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Group 1 elements
In the elemental state
Soft
 Silvery metal
 High melting points
 Extremely reactive therefore are not found in
elemental state in nature
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React violently with water to produce hydrogen
gas
Alkaline – Earth Metals
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Group 2 elements
Outer most s orbital is full
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Do not exhibit stability (outer p orbital is empty)
Properties
Harder, denser than group 1
 Higher melting points than group 1
 Not as reactive but too reactive to be found in
nature in elemental form
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Burning Mg
Hydrogen & Helium
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H has same valence electrons as group 1 but
does not share any other properties
He share same electron configuration (valence e) as group 2 but does not share same properties
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Placed with group 18 because it is very stable
d-block elements
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Transition elements
Beginning filling the 3d orbitals
 Good conductors of electricity
 High luster
 Less reactive than s-block elements
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Can be found in elemental form
Exceptions in the d-block
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The following elements have odd configurations
Cr: [Ar]4s13d5
 Cu: [Ar]4s13d10
 Ag: [Kr]5s14d10
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More stable with half filled s & d orbitals or full
d orbital
Exceptions follow throughout the d element
similar to Chromium and Copper
p-block elements
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All elements in p block have a full s orbital
Properties
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Contain all non metals except H & He
Contain all metalloids (exhibit properties of both metals and
non metals)
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Have semi conducting properties
Contains 6 metals
Elements in s & p block make up the representative
elements
Halogens
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Group 7A/17
Most reactive non metals (Fluorine is most reactive)
 Will bond with a metal to form a salt
 F & Cl are gases at room temp
 Br is a liquid at room temp
 I & At are solids at room temp
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Periodic Trends
Octet Rule
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Atoms will gain, lose, or share electrons in order
to have eight (8) valence electrons.
3 or less valence electrons – atom likely to lose
electrons
 6 or more valence electrons – atoms likely to gain
electrons
 4 or 5 valence electrons – atoms likely to share
electrons
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Periodic Trends
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Properties of the elements change in a
predictable manner across a period and down a
group
Atomic Radius
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The half distance
between nuclei of
identical atoms that are
chemically bonded
together
Atomic Radius
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Tends to decrease as you go across a period
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Increase nuclear charge pulls electrons closer to the
nucleus (decreasing radius) Zeff
Tends to increase as you go down a group
New electrons are placed in higher energy levels
 Shielding: core electrons shield outer electrons from
pull from nucleus
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Ionic Radius
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Ions – atom or bonded group of atoms that has
a positive or negative charge due to a loss/gain
of electrons
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Positive charge  lost electrons
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Smaller ionic radius compared to anions
Negative charge  gained electrons
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Larger ionic radius compared to cations
Ionic Radius
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Tends to decrease across a period
Tends to increase down a group
Ionic Radius vs. Atomic Radius
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Metals - the atomic radius of a metal is generally larger
than the ionic radius of the same element.
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Why? Generally, metals loose electrons to achieve the octet. This
creates a larger positive charge in the nucleus than the negative
charge in the electron cloud, causing the electron cloud to be
drawn a little closer to the nucleus as an ion.
Ionic Radius vs. Atomic Radius
cont.
Non-metals - the atomic radius of a non-metal is
generally smaller than the ionic radius of the same
element.
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Why? Generally, non-metals loose electrons to
achieve the octet. This creates a larger negative
charge in the electron cloud than positive charge in
the nucleus, causing the electron cloud to 'puff
out' a little bit as an ion.
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Ionic Radius vs. Atomic Radius
Ionization Energy
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Energy required to remove an electron from a
gaseous atom (J)
If an atom has a high ionization energy not likely
to form a positive ion
Tends to increase across a period
Tends to decrease down a group
st
1
Ionization Energy
Electronegativity
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Relative ability of an atom to attract electrons in
a chemical bond
Numerical value of 3.98 Paulings or less
Fluorine is the most electronegative
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atoms positioned closer to F have higher
electronegativies
Tends to increase across period
Tends to decrease down a group
Electronegativity
Reactivity
Reactivity refers to how likely or
vigorously an atom is to react with other
substances.
This is usually determined by two things:
1) How easily electrons can be
removed (ionization energy) from
an atom
2) or how badly an atom wants to
take other atom's electrons
(electronegativity)
The transfer/interaction of
electrons is the basis of chemical
reactions.
Reactivity of Metals
Period - reactivity decreases as you go from
left to right across a period.
Group - reactivity increases as you go down a
group
Why? The farther to the left and down the
periodic chart you go, the easier it is for
electrons to be given or taken away, resulting in
higher reactivity.
Reactivity of Non-Metals
Period - reactivity increases as you go from
the left to the right across a period.
Group - reactivity decreases as you go down
the group.
Why? The farther right and up you go on the
periodic table, the higher the
electronegativity, resulting in a more
vigorous exchange of electron.
Electron Affinity
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Change in energy that occurs when a neutral
atom acquires an electron
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Measured with a negative value
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Hopefully atom is become more stable by acquiring
an elcectron
The more negative the value the easier it is to
acquire an electron
Tends to become more negative across a period
Tends to become more positive down a group
Electron Affinity
Summary of Periodic Trends