1. Ch. 5 - The Periodic Table
I. History (p )

2. A. Mendeleev Dmitri Mendeleev (1869, Russian)
Organized elements by increasing atomic mass.
Elements with similar properties were grouped together.
There were some discrepancies.
C. Johannesson

3. A. Mendeleev Dmitri Mendeleev (1869, Russian)
Predicted properties of undiscovered elements.

4. B. Moseley Henry Mosely (1913, British)
Organized elements by increasing atomic number.
Resolved discrepancies in Mendeleev’s arrangement.
C. Johannesson

5. II. Organization of the Elements
Ch. 5 - The Periodic Table
II. Organization of the Elements

6. A. Metallic Character
Metals
Nonmetals
Metalloids

7. B. Blocks Main Group Elements Transition Metals
Inner Transition Metals

8. C. Periods

9. D. Families
Similar valence e- within a group result in similar chemical properties
C. Johannesson

10. D. Families Alkali Metals Alkaline Earth Metals Transition Metals
Halogens
Noble Gases
C. Johannesson

11. III. Periodic Trends (p. 140 - 154)
Ch. 5 - The Periodic Table
III. Periodic Trends (p )

12. A. Periodic Law
When elements are arranged in order of increasing atomic #, elements with similar properties appear at regular intervals.

13. B. Atomic Radius Atomic Radius size of the atom K Na Li Ar Ne
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Atomic Radius
size of the atom Li Ar Ne K Na

14. B. Atomic Radius Atomic Radius Increases to the LEFT and DOWN
C. Johannesson

15. B. Atomic Radius Why larger going down?
Higher energy levels have larger orbitals
Shielding - core e- block the attraction between the nucleus and the valence e-
Why smaller to the right?
Increased nuclear charge without additional shielding pulls e- in tighter
C. Johannesson

16. C. Ionic Radius Ionic Radius
The ionic radius of an element is its share of the distance between adjacent ions in an ionic solid

17. C. Ionic Radius Ionic Radius Cations (+) lose e- smaller Anions (–)
gain e-
larger
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C. Johannesson

18. C. Ionic Radius

19. D. Electronegativity
Electronegativity: affinity for gaining electrons from other elements.

20. D. Electronegativity

21. D. Electronegativity Why increase to the right? Closer to noble gas.
Why decrease going down?
Shielding

22. E. Ionization Energy First Ionization Energy
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First Ionization Energy
Energy required to remove one e- from a neutral atom. K Na Li Ar Ne He
C. Johannesson

23. E. Ionization Energy First Ionization Energy
Increases UP and to the RIGHT
C. Johannesson

24. E. Ionization Energy Why opposite of atomic radius?
In small atoms, e- are close to the nucleus where the attraction is stronger
Why small jumps within each group?
Stable e- configurations don’t want to lose e-
C. Johannesson

25. E. Ionization Energy Successive Ionization Energies
Large jump in I.E. occurs when a CORE e- is removed.
Mg 1st I.E kJ
2nd I.E. 1,445 kJ
Core e- 3rd I.E. 7,730 kJ
C. Johannesson

26. E. Ionization Energy Successive Ionization Energies
Large jump in I.E. occurs when a CORE e- is removed.
Al 1st I.E kJ
2nd I.E. 1,815 kJ
3rd I.E. 2,740 kJ
Core e- 4th I.E. 11,600 kJ
C. Johannesson

27. F. Electron Affinity
Electron Affinity - The amount of energy required, or given off, when a neutral atom gains an electron.
In a sense, it is the opposite of IE, since IE represents the energy change when an atom loses an electron.
These values are usually negative (indicating energy is given off.)

28. F. Electron Affinity
pg. 147 Modern Chemistry

29. F. Electron Affinity
Why increase energy released (become more negative) to the right within blocks?
Stability associated with full (and half full) valence shells.
Group trends not regular due to competing affects.
nuclear charge
Shielding (dominates)

30. G. Chemical Reactivity
Chemical Reactivity refers to how likely or vigorously an atom is to react with other substances, usually determined by how easily electrons can be removed (ionization E) and how attracted they are to other atom’s electrons (electronegativity).

31. G. Chemical Reactivity
Metals
Nommetals
increasing

32. H. Acid/Base Characteristics of Oxides

33. F. Melting/Boiling Point
Highest in the middle of a period.
C. Johannesson