1. Unit 3 – Periodic Table of Elements

2. 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

3. Unit 3 Key Terms (cont.)
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

4. 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

5. 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

6. Number of electrons (2n2)
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
Number of electrons (2n2) 1
2(12) = 2 2
2(22)= 8 3
2(32)= 18 4
2(42)= 32 n
2n2

7. 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
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

8. Total # of electrons per orbital type
Orbital Types
Orbital Type
General Shape
Orbital
Sublevels
# of electrons
per sublevel
Total # of electrons per orbital type s
Spherical 1 2 p
Dumbbell 3 6 d
Clover leaf 5 10 f
unknown 7 14

9. Electron Configuration
Energy Level
Orbital Type
Orbital
Sublevel
# of orbitals per energy level (n2)
# of electrons per orbital type
# of electrons per energy level (2n2) 1 s 2 p 3 4 6 8 d 5 9 10 18 f 7 16 14 32

10. 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 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p

11. Samples of e- Configuration
Element Electron Configuration
H 1s1
He 1s2
Li 1s2 2s1
C 1s2 2s2 2p2
K s2 2s2 2p6 3s2 3p6 4s1
V 1s2 2s2 2p6 3s2 3p6 4s2 3d3
Br 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5 (Note the overlap)
Pb 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p2

12. 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

13. 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

14. Valence Electrons
Shortcut to finding valence electrons for main group elements
Family 1A (1) valence electron
Family 2A (2) 2 valence electrons
Family 3A (13) 3 valence electrons
Family 4A (14) 4 valence electrons
Family 5A (15) 5 valence electrons
Family 6A (16) 6 valence electrons
Family 7A (17) 7 valence electrons
Family 8A (18) 8 valence electrons
Family 3-12 have multiple possibilities and shortcuts do not work

15. 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 Magnesium Chlorine Neon
Electron Dot Notation:
  • • •• ••
Na Mg : Cl : : Ne :
• • ••
Oxidation Numbers:

16. 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”

17. 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

18. 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

19. Mendeleev’s Periodic Table

20. 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

21. 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

22. Moseley’s Periodic Table
Image courtesy of

23. Development of Modern Periodic Table of Elements (cont.)
Glenn Seaborg ( United States )
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

24. Seaborg’s Periodic Table

25. Periodic Table of Elements 2012
Image used courtesy of

26. Dynamic Periodic Table
Courtesy of ptable.com
Image courtesy

27. 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

28. Types of Elements Metals Non-Metals Metalloids
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)

29. 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

30. 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

31. 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)

32. Periodic Trends