Chapter 6 The Periodic Table and Periodic Law

1. Chapter 6 The Periodic Table and Periodic Law The elements, which make up all living and non-living matter, fit into a orderly table. When interpreted properly, the table describes much of the elements physical and chemical properties.

2. What is the Periodic Law and how was it formulated? • Demitri Mendeleev is known as the father of the periodic table • He arranged the elements in families (groups) and periods (rows,series) according to atomic mass and properties

3. Mendeleev noted that the chemical and physical values for elemental properties would either be high or low depending upon the group under observation. • He proposed the first Periodic Law "The properties of the elements are a periodic function of their atomic masses" • Left blanks in his table for undiscovered elements

4. Moseley’s Modern-Periodic Law • There was some inconsistencies with Mendeleev’s table • In the early 1900’s Moseley was able to experimentally determine the atomic number of all known elements

5. Moseley then proceeded to rearrange the elements according to increasing atomic numbers. • New/Modern Periodic law states that the properties of elements are a periodic function of their atomic number

6. The Modern Periodic Table • Glenn T. Seaborg won the Nobel Prize for his work in nuclear chemistry • In 1944, formulated the “actinide concept” of heavy element electronic structure. This concept predicted that the fourteen actinides

7. Some Characteristics of Groups

8. Group 1 (IA) - Alkali Metals ( metal characteristics - shiny, malleable, ductile, good conductors) • Very active metals - activity increases as you go down group • All have one valence electron - form +1 cations by losing an electron • react violently with water

9. Group 2 (IIA) - Alkaline Earth Metals • activity increases as you move down the column not as reactive as alkali metals • Ca, Sr, and Ba react violently when they come into contact with water • All have two valence electrons • Form +2 cations by losing 2 electrons

10. Group 17 (VIIA) - Halogens • All gain one electron to form anions with a charge of -1. • All are nonmetals except for At which is a semimetal • All are diatomic in their elemental form

11. Group 18 (VIIIA) - Noble (Rare) Gases • Mistakenly labeled as "inert gases" until about 30 years ago because it was thought that these gases did not react with anything. • Noble gases have filled valence (outermost) shells.

12. Periodic Trends • Atomic Radii 1) As you move down a group, atomic radius increases. 2) WHY? - The number of energy levels increases as you move down a group as the number of electrons increases.

13. As you move across a period, atomic radius decreases. • WHY? - As you go across a period, electrons are added to the same energy level. At the same time, protons are being added to the nucleus. The concentration of more protons in the nucleus creates a "higher effective nuclear charge."

14. First Ionization Energy • Definition: The energy required to remove the outermost (highest energy) electron from a neutral atom in its ground state. 1) As you move down a group, first ionization energy decreases. 2) WHY? Electrons are further from the nucleus and thus easier to remove the outermost one - + shielding effect of other electrons

15. 3) As you move across a period, first ionization energy increases. 4) WHY? - As you move across a period, the atomic radius decreases, the attraction from the positive nucleus gets larger

16. Electronegativity • Definition: The ability of an element to attract electrons in a chemical bond 1) As you move down a group, electronegativaty decreases 2) As you move across a period, it increases.