THE PERIODIC TABLE Part I – The History

1. THE PERIODIC TABLEPart I – The History

2. First Attempts to Organize the Elements: Dobereiner (1826) - recognized that there were several groups of 3 elements with similar properties, “TRIADS”. Newlands (1864) – in organizing the elements by the atomic masses, noticed that every 8th element seemed to be repeat, “the Law of Octaves.”

3. Dmitri Mendeleev The Father of the Modern Periodic Table In 1869, organized the elements known at the time (66) according to: - Increasing Atomic Mass - Physical and chemical properties

4. Mendeleev’s Periodic Law The physical and chemical properties of the elements are periodic functions of their atomic masses

5. Mendeleev left “holes” in his Periodic Table Since only 66 elements were know when Mendeleev came up with his table, he left spaces where no known elements matched the properties of the elements in the table.

6. Mendeleev predicted the properties of these undiscovered elements He predicted the properties of Gallium (which he called eka-aluminum) and Germanium (which he called eka-silicon) amongst several others.

7. Comparison of the predicted properties compared to the actual properties.

8. But there was a problem! The mass of iodine was actually greater than tellurium. This was a problem because the chemical and physical properties of Iodine were consistent with bromine and chlorine, not Selinium and Sulfur. Mendeleev opted to align Iodine with bromine and chlorine, despite the fact that it seemed to violate his Periodic Law.

9. Problem Solved! The discrepancy involving Iodine and Tellurium was resolved when Henry Mosely discovered the concept of Atomic Number in 1913. Mosely found that although Tellurium was heavier than Iodine, its atomic number (52) was less than Iodine (53).

10. The Modern Periodic Law The physical and chemical properties of the elements are periodic functions of their atomic numbers.

11. THE PERIODIC TABLEPart II – The Structure

12. Organization of the Periodic TablePeriods • Periods - The horizontal rows in the Periodic Table. Elements in the same period have different properties. • Examples: H and He are in the 1st Period, Li through Ne make up the 2nd Period.

13. Organization of the Periodic TableGroups (or Families) • Groups - The vertical columns in the Periodic Table. Elements in the same group have similar properties.

14. Alkali Metals Group 1 Li, Na, K, Rb, Cs, Fr Very reactive, not found uncombined in nature

15. Alkaline Earth Elements Group 2 Be, Mg, Ca, Sr, Ba, Ra Very reactive, not found uncombined in nature.

16. Halogens Group 17 (VIIA) F, Cl, Br, I, As Very reactive, not found uncombined in nature

17. The Noble Gases Group 18 He, Ne, Ar, Kr, Xe, Rn Very unreactive. Do not normally react with other elements to form compounds.

18. The Transition Metals Groups 3 through 12 Form colored compounds

19. The Rare Earth Elements The inner transition elements Found between Groups 3 and 4 Also know as the Lanthanide and Actinide Series.

20. Metals Most elements in the Periodic Table are metals. They have luster (shiny), and are malleableand good conductors of electricity. They tend to lose electrons in chemical reactions. Metals are located on the left side of the Periodic Table.

21. Non-Metals They are brittle and are poor conductors of electricity. They tend to gain electrons in chemical reactions. Non-Metals are located on the right side of the Periodic Table.

22. Metalloids (Semi-metals) They have some properties associated with metals and some properties associated with non-metals. One of their most important properties is that they are semi-conductors (used in electronic devices like cell phones and computers. Metalloids are located adjacent to the “staircase of chemistry.”

23. Homework Assignment • Using the blank Periodic Table provided. • Draw electron dot diagrams for each element in Groups 1, 2,13,14,15,16,17 and 18 • On the table, indicate the following groups using different colors (with a key) : • alkali metals • alkaline earth elements • halogens • noble gases • transition elements • rare earth elements.

24. The Periodic Table and Electron Structure Although the Periodic Table originally was organized based on properties, in doing so, it was also organized by electron structure: **** Elements in the same group have the same number of valence electrons. **** Elements in the same period have their valence electrons in the same energy level.

25. Ionization Energies Ionization energy is defined as the amount of energy required to remove an electron from an atom. To help you understand the ionization energy trends in the Periodic Table, you will need to graph the 1st Ionization Energies for the first 18 elements (y axis) vs. their atomic numbers (x axis)

26. On the graph paper provided, prepare two copies of the graph of ionization energies (from Table S) vs. Atomic Number for elements 1-18. By Period By Group

27. In the Graph labeled “By Period”, connect the data points for elements in the same Period. What happens to ionization energy as you go across a Period? ****In general, Ionization energy increases as you go across a period.

28. In the Graph labeled “By Group”, connect the data points for elements in the same Group. What happens to ionization energy as you go down a Group? ****In general, Ionization energy decreases as you go down a group.

29. Atomic Radii Decreases Across a period? Down a group? Increases

30. Electronegativity Measure of an atoms ability to attract electrons in a bond between two atoms. Across a period: Increases Down a group: Decreases

31. Metallic Character (Properties) Refers to an atoms ability to lose electrons. The more easily an atom loses electrons the more metallic it is.

32. Summary of the trends Visit this web page for an explanation of the trends.