1. Creativity at its best!  Consider the periodic table. People could have invented various tables in the 17 th century but Mendeleev created the periodic table of elements. Ponder on your interests and hobbies in life. If you could create your own periodic table, what would it be in relations to? How would it be organized? Brainstorm this and write about it.

2. Development of the Periodic Table Sizes of Atoms and Ions Ionization Energy Electron Configurations

3. E.Q.: Can I compare and contrast trends in the chemical and physical properties of elements and their placement on the Periodic Table (SC4b)?

4.  Metals (+)  Metalloids (+)  Nonmetals (-)  Developed by Dimitri Mendeleev and Julius Meyer.  Mendeleev is given most credit.  Among the first discovered were iron (Fe), copper (Cu), gold (Au), and silver (Ag).  Initially arranged in order of mass number.

5.  Henry Moseley later developed the concept of atomic numbers.  The table was then arranged in order of increasing atomic number.  Majority of elements exist in nature only as compounds.

6. 6 Except for H, elements left of the zigzag line are metals. To the right of the line we find nonmetals, including the noble gases. Some elements adjacent to the line are called metalloids.

7.  Because e - are negative, they’re attracted to nuclei, which are _____________.  Coulomb’s Law – The strength of attraction between two electric charges depend on...  Thus, the force of attraction between electrons and nuclei depend on...

8.  In many – e - atoms, each e - is attracted to the nucleus and repelled by...  The charge resulting from positive charge of the nucleus and the shielding effects of e - is called the effective nuclear charge.  The net electric charge of the nucleus after the shielding effects of its surrounding e -

9.  Effective nuclear charge increases as we move across any period of the table.  Effective nuclear charge increases as we move down a group...HOWEVER, not as much as moving across a period.

10.  Indicate which would you expect to experience a greater effective nuclear charge: a Neon atom or a sodium atom.

11.  Creation of the periodic table...  Development of the modern periodic table...  Effective nuclear charge...

12. Study your notes!

13.  Fill in the diagram below and tell me what this diagram is demonstrating.  s: _____  p: _____ _____ _____  d: _____ _____ _____ _____ _____  f: _____ _____ _____ _____ _____ _____ ____

14. Development of the Periodic Table Sizes of Atoms and Ions Ionization Energy Electron Configurations

15. E.Q.: How do I use the periodic table to predict periodic trends including atomic radii, ionic radii, ionization energy, and electronegativity (SC4a)?

16.  Remember this concept?  Explain how effective nuclear charge increases and decreases.  Which would you expect to experience a greater effective nuclear charge, helium or magnesium?

17.  One of the most important properties of atoms/ions – SIZE!!!  What determines size of atoms/ions?  Atomic radius (plural: radii)  Distance from nucleus to valence electrons.  Measured in Angstroms (1.0 Å = 10 -10 m)  Decreases across table; increases down table.

18.  Look at the trends...

19.  Loss/gain of an electron...  s: _____  p: _____ _____ _____  d: _____ _____ _____ _____ _____  f: _____ _____ _____ _____ _____ _____ ____

20.  What determines size of ions?  Formed when atoms lose/gain e -  Less e - are pulled more closely to nucleus.  More e - - larger radius, less e - - smaller radius.  Positive ions are smaller than their parent atoms!  Negative ions are larger than their parent atoms!

22.  Atomic radii________ as you move from left to right across a period.  Atomic radii ________ as you move down a group.  Ionic radii __________ as you gain e-.  Ionic radii __________ as you lose e-.

23.  Arrange these atoms and ions in order from smallest to largest: B, F, and O.

24.  Sizes of Atoms and Ions...  Periodic Trends in Atomic Radii...  Periodic Trends in Ionic Radii...

25. Study Your Notes

26.  Based on what you know, opine why atoms have different sizes.  Observe the equation below. Opine what’s going on in the chemical equation:

27. Development of the Periodic Table Sizes of Atoms and Ions Ionization Energy Electron Configurations

28. E.Q.: How do I use the periodic table to predict periodic trends including atomic radii, ionic radii, ionization energy, and electronegativity (SC4a)?

29.  Minimum energy required to remove an e - from an atom.  First ionization energy (I 1 ) – Minimum energy required to remove the first e - from an atom.  Second ionization energy (I 2 ) – Minimum energy required to remove the second e - from an atom. So on and so fourth!

30.  The greater the ionization energy, the more difficult it is to remove an electron!

31.  Ionization energy values increase as successive electrons are removed.  I 1 < I 2 < I 3

35.  Referring to the periodic table, rank the following atoms in order of smallest to largest ionization energy: Ne, Na, P, Ar, and K.

36.  Ionization Energy...  Periodic trends in Ionization Energy...

37. Study your notes!

38.  Fill in the electron orbital diagram below  s: ____  p: ____ ____ ____  d: ____ ____ ____ ____ ____  f: ____ ____ ____ ____ ____ ____ ____

39. Development of the Periodic Table Sizes of Atoms and Ions Ionization Energy Electron Configurations

40. E.Q.: How do I use the periodic table to predict periodic trends in electron configuration?

41.  Electron configuration - Distribution of electrons among orbitals in atoms.  Represented in two ways: 1. spdf notation 2. Orbital box notation

42.  Coefficient – Shell  s,p,d,f – designates orbital.  Superscript – How many e - are in that orbital.

43. Represents electrons within orbitals and orbitals within shells. The arrows’ directions represent electron spins; opposite spins are paired. N:

44.  e - ordinarily occupy orbitals of the lowest energy available.  No two e - in the same atom may have all four quantum numbers alike.  Pauli exclusion principle: one atomic orbital can accommodate only two e -, and these e - must have opposing spins.  e - in half-filled orbitals have parallel spins (same direction).

45. The electron configuration of Si ends with 3s 2 3p 2 The electron configuration of Rh ends with 5s 2 4d 7

46.  4d  5p  5s  3d  4p  4s  3p  3s  2p  2s  1s

47.  Try study guide questions (#5).

48.  Sc: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 1  Ti:

49.  Illustrate the electron configuration for hydrogen and lithium.  Indicate the element with the following electron configuration: 1s 2 2s 2 2p 4

50.  Electron configuration...

51. Study Your Notes!