1. Chapter 6 The Periodic Table and Periodic Law
Section1: History
Section 2: Classification of the Elements
Section 3: Periodic Law

3. 1) John Newlands…
2) Meyer…
3) Mendeleev:
He left blank spaces in his periodic table where he thought undiscovered elements had to go and predicted their properties from periodic trends.
Organized his periodic table by atomic mass
4) Moseley
Revised periodic table, sorting it by atomic numbers, demonstrating periodic law

4. The Modern Periodic Table
Periods (  rows)
Groups (columns)
Representative elements (main groups, long columns)
Transition elements (all metals, in the middle)
Inner transition elements (metals underneath)
Metals/nonmetals/metalloids
Alkali metals (groups 1)
Alkaline earth metals (groups 2)
Halogenes (group 17)
Noble gases (groups 18, inert)

5. Periodic Table Metals blue Nonmetals in upper right corner, brown,
Metalloids along the stairs separating metals and nonmetals
Letter A: Main group or Representative elements (long columns)
Letter B: Transition metals and inner transition metals

6. Section 2 Classification of the elements
Valence electrons
s, p, d, f blocks

7. Section 3
Periodic trends
Atomic radius
Ionic radius
Ionization energy
Electronegativity

8. Atomic size is a periodic trend influenced by electron configuration
For metals, atomic radius is half the distance between adjacent nuclei in a crystal of the element.
For elements that occur as molecules, it is half the distance between nuclei of identical atoms.

10. Shielding Effect
The decrease in the attraction between outer electrons and the nucleus due to the presents of other electrons between them
Shielding effect increases down a group in the periodic table as more and more energy levels are added, but is constant along a period (left to right).

11. There is a general decrease in atomic radius from left to right, caused by increasing positive charge in the nucleus.
Valence electrons are not shielded from the increasing nuclear charge because no additional electrons (energy level n)come between the nucleus and the valence electrons.
Atomic radius generally increases as you move down a group.
The outermost orbital size increases down a group, making the atom larger.

14. Ionic radii
Ions change their size based on the new ratio of electrons to protons. Less electrons feel the positive charge stronger, more electrons feel it less. Furthermore, cations often lose a complete energy level, decreasing the size significantly.

16. The ionic radii of positive ions generally decrease from left to right.
The ionic radii of negative ions generally decrease from left to right, beginning with group 15 or 16.

17. [Na  Na+ + e- ] formation of a positive ion.
Ionization Energy
The energy required to overcome the attraction of the nuclear charge and remove an electron from a gaseous atom is called the ionization energy.
[Na  Na+ + e- ] formation of a positive ion.

18. Removing the second electron requires more energy, and is called the second ionization energy.
Each successive ionization requires more energy, but it is not a steady increase.

19. The octet rule states that atoms tend to gain, lose or share electrons in order to acquire a full set of eight valence electrons.
The octet rule is useful for predicting what types of ions an element is likely to form.

20. The electronegativity of an element indicates its relative ability to attract electrons in a chemical bond.

22. The lowest ionization energy is the ____.
A. first
B. second
C. third
D. fourth