1. Unit 3 – Periodic Table of Elements

2. Early Development of the Periodic Table of Elements
Antoine Lavoisier
John Dalton
Jacob Berzelius
Johann Dobereiner
John Newlands

3. Development of Modern Periodic Table of Elements
Dmitri Mendeleev (Russia 1864)
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
Henri Becquerel
Frederick Soddy
J.J. Thomson

4. Mendeleev’s Periodic Table

5. 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

6. Moseley’s Periodic Table
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7. 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

8. Seaborg’s Periodic Table

9. Periodic Table of Elements 2012
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10. Dynamic Periodic Table
Courtesy of ptable.com
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11. 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

12. 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)

13. 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

14. 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

15. 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)

16. Periodic Trends