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Unit 3 – Periodic Table
of Elements
Unit 3 Key Terms
• Energy Level – discrete regions of space around the
nucleus in the electron cloud where electrons can reside
• Lewis dot structure -A model that uses electron-dot
structures to show how electrons are arranged in
molecules. Pairs of dots or lines represent bonding pairs
• Noble gas configuration -An electron structure of an
atom or ion in which the outer electron shell contains
eight electrons, corresponding to the electron
configuration of a noble gas, such as neon or argon
• Orbital notation (diagram) -A way to show how many
electrons are in an orbital for a given element. They can
either be shown with arrows or circles
Unit 3 Key Terms (cont.)
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Atomic Mass – physical property indicating the mass of an atom
Alkali Metals – extremely reactive metals found in Family 1A
Alkaline Earth Metal – very reactive metals found in Family 2A
Atomic Radius – one half the distance between two adjacent
atoms of the same element
Electronegativity – the ability of an element to attract electrons
from a neighboring atom
Ionization Energy – the energy required to remove an electron
from the electron cloud of an element
Halogens – extremely reactive non-metals found in Family 7A
Ion – charged atom resulting from the loss or gain of electrons
Ionic Radius – the resulting atomic radius found in an element
when it has either lost or gained electrons to become an ion
Unit 3 Key Terms (cont.)
• Mendeleev – Russian chemist that placed elements into a
Periodic Table with periods based on atomic mass and families
(groups) based on similar chemical and physical properties
• Mosley – English physicist that used X-ray spectroscopy to find
the Atomic Number of elements and place elements into a
Periodic Table with periods based on atomic number and families
(groups) based on similar chemical and physical properties
• Noble Gases – stable, non-reactive non-metals in Family 8A
• Periodic Trends – repeating patterns of chemical and physical
properties of elements within the periodic table that correspond
to the Law of Periodicity
• Reactivity – chemical property describing the ability and speed
with which elements react
• Transition Metals – metals in the d-block, generally with multiple
valence states; most of the commonly known metals are in this
category of metals
Quantum Mechanical Model of
Atomic Structure
• 1900: Max Planck – Develops law correlating energy to
frequency of light
• 1905: Albert Einstein – Postulates dual nature of light as both
energy and particles
• 1924: Louis de Broglie – Applies dual nature of light to all
matter
• 1927: Werner Heisenberg – Develops Uncertainty Principle
stating that it is impossible to observe both the location and
momentum of an electron simultaneously
• 1933: Erwin Schrodinger – Refines the use of the equation
named after him to develop the concept of electron orbitals to
replace the planetary motion of the electron
Energy Levels
• Energy levels correspond to the energy of individual
electrons. Each energy level has a discrete numerical
value.
• Different energy levels correspond to different numbers
of electrons using the formula 2n2 where “n” is the
energy level
Energy Level
1
2
3
4
n
Number of electrons (2n2)
2(12) = 2
2(22)= 8
2(32)= 18
2(42)= 32
2n2
Orbitals
 Impossible to determine the location of any single electron
 Orbitals are the regions of space in which electrons can most
probably be found
 Four types of orbitals
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s – spherically shaped
p – dumbbell shaped
d – cloverleaf shaped
f – shape has not been determined
 Each additional energy level incorporates one additional orbital type
 Each type of orbital can only hold a specific number of electrons
Orbital Types
Orbital
Type
General
Shape
Orbital
Sublevels
1
# of
electrons
per
sublevel
2
Total # of
electrons
per orbital
type
2
s
Spherical
p
Dumbbell
3
2
6
d
Clover leaf
5
2
10
f
unknown
7
2
14
Electron Configuration
Energy
Level
Orbital
Type
Orbital
Sublevel
1
s
1
s
p
s
p
d
s
p
d
f
1
3
1
3
5
1
3
5
7
2
3
4
# of
# of
# of
orbitals
electrons electrons
per energy per orbital per energy
level (n2)
type
level (2n2)
1
4
9
16
2
2
6
2
6
10
2
6
10
14
2
8
18
32
Electron Configuration
Notation
• Find the element on the periodic table
• Follow through each element block in order by stating the
energy level, the orbital type, and the number of electrons per
orbital type until you arrive at the element.
1s
2s
3s
4s
5s
6s
7s
4f
5f
3d
4d
5d
6d
2p
3p
4p
5p
6p
7p
Samples of e- Configuration
• Element Electron Configuration
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H
He
Li
C
K
V
Br
Pb
1s1
1s2
1s2 2s1
1s2 2s2 2p2
1s2 2s2 2p6 3s2 3p6 4s1
1s2 2s2 2p6 3s2 3p6 4s2 3d3
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5 (Note the overlap)
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p2
Noble Gas Electron Configuration Notation
 Find element on the Periodic Table of Elements
• Example: Pb for Lead
 Move backward to the Noble Gas immediately preceding the
element
 Example: Xenon
 Write symbol of the Nobel Gas in brackets
 Example: [Xe]
 Continue writing Electron Configuration Notation from the
Noble Gas
 Example: [Xe] 6s2 4f14 5d10 6p2
Valence Electrons
• The electrons in the highest (outermost) s and p orbitals of an
atom
• The electrons available to be transferred or shared to create
chemical bonds to form compounds
• Often found in incompletely filled energy levels
Valence Electrons
 Shortcut to finding valence electrons for main group
elements
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Family 1A (1)
Family 2A (2)
Family 3A (13)
Family 4A (14)
Family 5A (15)
Family 6A (16)
Family 7A (17)
Family 8A (18)
1 valence electron
2 valence electrons
3 valence electrons
4 valence electrons
5 valence electrons
6 valence electrons
7 valence electrons
8 valence electrons
 Family 3-12 have multiple possibilities and shortcuts do not
work
Electron Dot Notation
 Electron configuration notation using only the valence electrons of an
atom.
 The valence electrons are indicated by dots placed around the element’s
symbol.
 Used to represent up to eight valence electrons for an atom. One dot is
placed on each side before a second dot is placed on any side.
Valance Electrons:
Sodium
1
Electron Dot Notation:
•
Na
Oxidation Numbers:
+1
Magnesium
2
Chlorine
7
Neon
8
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Mg
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: Cl :
: Ne :
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••
+2
-1
0
Early Development of the
Periodic Table of Elements
• Antoine Lavoisier (France 1778)
• Produced the first extensive list of elements showing 33 elements
• Separated metals from non-metals
• John Dalton (England 1803)
• Developed postulates of atomic theory with a list of elements and
symbols
• Jacob Berzelius (Sweden 1828)
• Systematized letters to symbolize elements
• Provided a table of atomic weights
• Johann Dobereiner (German 1828)
• Discovered patterns between elements in groups of 3
• Elements in triads formed similar chemical compounds
• Published list of these groups called “triads”
Early Development of the
Periodic Table of Elements (cont.)
• John Newlands (England 1864)
• First person to devise a periodic table of elements
• Expanded the concept of triads into octaves. Elements were said
to exhibit similar chemical and physical properties to the eighth
element following it in the table (Law of Octaves)
• His table did not include all of the known elements
Development of Modern
Periodic Table of Elements
• Dmitri Mendeleev (Russia 1869)
• Produced the first Periodic Table to arrange elements in periods
(rows) and families (columns) showing all 66 known elements
• Periods arranged elements in order of increasing atomic mass
• Families arranged by similar chemical and physical properties
• Method of arrangement left gaps for elements believed to exist and
not yet discovered.
• William Ramsay (England 1894)
• Discovered a new family of gases that resisted chemical reactions
• Noble gases added to the Periodic Table
• Henri Becquerel (France 1903)
• With Pierre and Marie Curie credited with discovering radioactivity
• Opened a new window on understanding atomic structure and
properties
Mendeleev’s Periodic Table
Development of Modern Periodic
Table of Elements (cont.)
• Frederick Soddy (England 1912)
• Suggested existence of isotopes
• Since elements could have multiple masses they could occupy
multiple positions on the Periodic Table arranged by mass
• J.J. Thomson (England 1913)
• Confirmed experimentally the existence of isotopes with Neon
• Since Neon is unreactive its multiple masses could not be explained
by unidentified chemical compounds
• Discovery of isotopes meant that the Mendeleev Periodic Table could
not be the final answer
Development of Modern Periodic
Table of Elements (cont.)
• Henry Moseley (England 1913)
• Demonstrated through x-ray spectroscopy that the characteristics of
the x-rays emitted by different atoms are incremental and can be
listed in numerical order (Atomic Number)
• Put forward the theory that chemical and physical properties are
periodic functions of this atomic number (Law of Periodicity)
• Refined Rutherford’s theory of the atomic structure indicating a
correlation between the positive charge of the nucleus and atomic
number
• Developed the basic structure of the Periodic Table used today.
• Rutherford (England 1917)
• Produced the first experimentally based nuclear reaction
transforming nitrogen into oxygen using alpha particles
• The other product of the reaction was a Hydrogen nucleus (Proton)
• Confirmed atomic number was equal to the number of protons
Moseley’s Periodic Table
Image courtesy of http://corrosion-doctors.org/Periodic/Periodic-Moseley.htm
Development of Modern Periodic
Table of Elements (cont.)
• Glenn Seaborg ( United States 1941-1944)
• Discovered discrepancies in Moseley’s table through the
identification of new elements while conducting research as part
of the Manhattan Project
• Created the lanthanide and actinide series referred to as
transuranium elements
• Discoveries disclosed at the end of World War II
• Continuing research
• Research laboratories use particle accelerators to identify new
elements
• Recent discoveries have completed the 7th period of the table
• Research is continuing to discover more new elements in an 8th
period
Seaborg’s Periodic Table
Periodic Table of Elements 2012
Image used courtesy of https://proteabio.com/resources/tools/Periodic+Table+of+Elements
Dynamic Periodic Table
• Courtesy of ptable.com
Image courtesy http://www.fanpop.com
Families of Particular Importance
 Family 1A (1) – Alkali Metals
 Soft metals and silver gray in color
 Extremely reactive – do not exist in elemental form in nature
 1 valence electron
 Family 2A (2) – Alkaline Earth Metals
 Soft metals and silver in color
 Very reactive – can exist in nature, but oxidize rapidly
 2 valence electrons
 Family 7A (17) – Halogens (non-metals)
 Very reactive
 Lighter halogens are gases at room temperature while heavier halogens are solids
at room temperature; Bromine is a liquid at room temperature
 7 valence electrons
 Family 8A (18) – Noble Gases (Non-metals)
 Generally non-reactive and do not form compounds
 Extremely Stable
 8 valence electrons
Types of Elements
 Metals
 Good conductors of electricity and heat
 Vast majority of the elements – Alkali metals, alkaline earth metals,
transition metals, post-transition metals, and inner transition metals
 Non-Metals
 Poor conductors of electricity and heat
 Includes the Nobel Gases, Halogens, and only a few of the lightest
elements – hydrogen, carbon, nitrogen, oxygen, phosphorus, sulfur, and
selenium
 Metalloids
 Many are semiconductors of electricity
 Exhibit properties of both metals and non-metals
 Only 7 elements are metalloids (some scientists include different ones
depending on perspective)
Periodic Law
Created by Henry Moseley
The chemical and physical properties of the elements are
periodic functions of their atomic numbers
Properties of the elements occur at repeated intervals called
periods (rows on Periodic Table)
This defines the property of periodicity
Periodic Trends
 Atomic Radius – half the distance between the nuclei of atoms of
the same material
 Decreases generally across periods – increased positive nuclear
charge (protons) pulls electrons in tighter to the nucleus
 Increases generally down families – increased number of energy
levels where electrons may reside
 Electronegativity – the measure of the ability of the nucleus of an
atom to attract electrons of a neighboring atom
 Increases generally across periods – increased positive nuclear charge
(protons) more strongly attracts electrons from neighboring atoms
 Decreases generally down families – increased number of energy
levels means the nucleus is less able to overcome the distance
between atom
Periodic Trends
 Ionization Energy – the energy required to remove an electron
from an atom
 Increases generally across periods – increased positive nuclear charge
(protons) pulls electrons in tighter to the nucleus making them harder
to remove
 Decreases generally down families – increased number of energy
levels where electrons may reside making them easier to remove
 Ionic Radius – as atoms gain and lose electrons, the radius of the
charged atom changes
 Increases when an atom accepts electrons– the more electrons there
are the greater the overall repulsive forces between the electrons
pushing them further apart
 Decreases when an atom loses electrons – the fewer the electrons the
greater the effectiveness of the nuclear charge (protons)
Periodic Trends