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Chapter 6 The Periodic Table and Periodic Law

Chapter 6 The Periodic Table and Periodic Law. Section 6.1. Development of the Modern Periodic Table. Objectives. Trace the development and Identify key features of the periodic table. History. 1790s, French scientist Lavoisier 23 elements 1800s: electricity and spectrometer.

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Chapter 6 The Periodic Table and Periodic Law

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  1. Chapter 6The Periodic Table and Periodic Law

  2. Section 6.1 Development of the Modern Periodic Table

  3. Objectives • Trace the development and Identify key features of the periodic table.

  4. History • 1790s, French scientist Lavoisier • 23 elements • 1800s: electricity and spectrometer

  5. John Newlands • Patterns of the periodic table • Law of octaves • Elements on the eights had similarities

  6. Meyer & Mendeleev • 1869-Meyer and Mendeleev: connection between atomic mass and properties • Mendeleev published it first • Left holes of undiscovered elements • Predicted properties of undiscovered elements- Sc, Ga, Ge

  7. Moseley • Problems with Mendeleev • Rows by increasing atomic mass • Moseley used atomic number instead

  8. The Periodic Law • There is a periodic repetition of chemical and physical properties of the element when they are arranged by increasing atomic number.

  9. The Modern Periodic Table • Horizontal rows- periods • Vertical columns- groups or families

  10. Groups 1A through 8A = Representative Elements • Wide range of chemical and physical properties • Groups 1B through 8B = Transition Elements

  11. Metals • shiny, smooth • solid room temperature • good conductors of heat and electricity

  12. Alkali Metals • Alkali Metals= 1A (excluding hydrogen) • highly reactive

  13. Alkaline Earth Metals • Alkaline Earth Metals= 2A • highly reactive (not as much as 1A)

  14. Transitions • Transition metals • Group B elements contained in the D block of the table • Inner transition metals • the lanthanide and actinide series • F block

  15. Nonmetals • Generally a gas or a brittle, dull-looking solids • Poor conductors • Halogens=7A • REALLY REACTIVE • Noble Gases= 8A • unreactive and stable (all valence electrons are filled)

  16. Metalloids • contain the physical and chemical properties of both metals and nonmetal

  17. Homework • Section 6.1 Assessment • 1-6 on page 158

  18. Reactivity Trends • Increases down the table for the metals • Decreases down the table for the nonmetals

  19. Section 6.2 Classification of the Elements

  20. Objectives • Explain why elements in the same group have similar properties • Identify the four blocks of the periodic table based on electron configuration

  21. Valance Electrons • The properties of each element in each group are similar because they have the same number of valence electrons

  22. Valence Electrons and Period Number • Energy level of valence electrons = period on the table • Example: Lithium: period 2, valence electron in 2nd energy level (1s22s1)

  23. Valence Electrons and Group Number • Group number (1A-8A) = number of valence electrons

  24. Putting it Together • Oxygen is in the 3rd period and in group 6A. • Oxygen’s valence electrons are located in the ____ energy level. • Oxygen has ___ valence electrons.

  25. S,P,D, and F Blocks • 4 different energy sublevels: s, p, d, and f • S block= 1A and 2A • holds max of 2 electrons • P block= 3A through 8A • max holds 6 electrons • S block must fill before P block can fill • Noble gases are stable because of filled S and P blocks

  26. S,P,D, and F Blocks Continued… • D block = transition metals • max of 10 electron • F block= inner transition metals • unpredictable manner of filling • max of 14 electrons

  27. Reactivity Trends • Increases down the table for the metals • Decreases down the table for the nonmetals

  28. Homework • Page 162, #10-15

  29. Periodic Trends: Atomic Radius For metals, atomic radius = ½ distance between adjacent nuclei in a crystal element For nonmetals, atomic radius = ½ distance between nuclei of identical atoms that are chemically bonded

  30. Periodic Trends: Atomic Radius DECREASES I N C R E A S E S

  31. Atomic Radius • The increase from top to bottom is due to adding electron shells. • The decrease from left to right is due to increased nuclear charge as you move to the right, which draws electrons closer to the nucleus.

  32. Periodic Trends: Ionization Energy • Ionization energy= energy required to remove an electron from a gaseous atom • Octet rule = atoms tend to gain lose or share electrons to acquire a full set of 8 valence electrons

  33. Periodic Trends: Ionization Energy INCREASES D E C R E A S E S

  34. Periodic Trends: Electronegativity • Electronegativity= relative ability of an atom to attract electrons in a chemical bond.

  35. Electronegativity • Arbitrary units called Paulings (after Linus Pauling) are used to express electronegativity. • Electronegativity increases from left to right across a period and from top to bottom down a group.

  36. Periodic Trends: Electronegativity INCREASES D E C R E A S E S

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