1. Chapter 14 - The Periodic Table and Periodic Law Objectives: –Identify different key features of the periodic table. –Explain why elements in a group have similar properties. –Relate the group and period trends seen in the periodic table to the electron configuration of atoms. Why this is important: –The periodic table is one of the most useful reference tools available in chemistry! Understanding its organization and interpreting its data will aid in understanding chemistry concepts. www.privatehand.com/flash/elements.html

2. Development of the Periodic Table In 2003, there were 118 elements known. The majority of the elements were discovered between 1735 and 1843. How do we organize all the different elements in a meaningful way that will allow us to make predictions about undiscovered elements?

3. 1869 Dmitri Mendeleev and Lothar Meyer separately arranged the elements in order of increasing atomic mass and into columns with similar properties. Mendeleev is given more credit than Meyer because he published his findings first, and he left spaces for elements that were not yet discovered. Some of the elements that he predicted were scandium, gallium, and germanium. –In 1871, Mendeleev noted that arsenic (As) properly belonged underneath phosphorus (P) and not silicon (Si), which left a missing element underneath Si. He predicted a number of properties for this element. In 1886 Germanium (Ge) was discovered. The properties of Ge matched Mendeleev’s predictions. http://www.chemistrydaily.com/chemistry/upload/a/a1/Dmendeleev.jpg http://www.chemistryexplained.com/images/chfa_03_img0535.jpg

4. Mendeleev’s table was not completely correct. Arranging elements by atomic mass caused some elements to be put in the wrong groups so that the properties did not exactly match up. 1913 English chemist Henry Moseley arranged elements in order of increasing atomic number. Problems with order of elements were solved, and there was a clear repeating pattern of properties of the elements in their groups. The PERIODIC LAW states there is a “periodic” repetition of chemical and physical properties of the elements when they are arranged by increasing atomic number. Periodic means: happening or reoccurring at regular intervals. (definition from Webster’s Dictionary) http://www.rsc.org/education/teachers/learnnet/periodictable/scientists /moseley.jpg

5. The Modern Table Boxes are arranged in order of increasing atomic # Elements are grouped into columns by similar properties Scientists keep adding elements that were discovered The final adjustment was when physicist Glenn Seaborg had the inner-transition elements pulled below the rest of the periodic table and into 2 separate rows (This occurred in the late 1940s) http://www.lbl.gov/Science-Articles/Research- Review/Magazine/1994/seaborgium-mag.html

6. Horizontal rows are called periods There are 7 periods 12345671234567

7. Vertical columns are called groups or families. Elements are placed in columns by similar properties. –b/c of the similar numbers of valence e - they contain!

8. 12 1314151617 18 3456789101112 The Different Groups of Elements 1 – 18 system used by all chemists A & B system is an older American System A elements are representative elements B elements are transition elements 1A2A 3A4A5A6A7A 8A 3B4B5B6B7B8B 1B2B

9. Representative or Main Group Elements Wide range of physical & chemical properties. The whole range of possible valence electrons (1 to 8) Also called s and p block elements Here are some important groups:

10. Group 1 (1A) contains the alkali metals (remember to NOT include hydrogen) Group 2 (2A) contains the alkaline earth metals

11. Alkali Metals These metals react w/ water to form alkaline (basic) solutions. Highly reactive metals that lose their 1 valence electron to form 1+ ions. Soft enough to be cut with a knife. They are stored in oil to prevent reactions with oxygen and water in the air.

12. Alkaline Earth Metals Most of these metals react with oxygen to form compounds called oxides (the alchemists called them “earths” because of this) and the oxides react w/ water to form alkaline (basic) solutions. Not as reactive (but do react easily) & harder than group 1 metals. They lose their 2 valence electrons to form 2+ ions.

13. Group 17 (7A) contains the halogens Group 18 (8A) contains the noble gases

14. Halogens Halogen means “salt formers” b/c they react with metals to form salts (ionic compounds). F & Cl are gases at room temp., Br is a liquid but it evaporates easily, and Iodine is a solid that sublimes easily. Astatine is radioactive w/ no known uses. They are the most reactive nonmetals! 7 valence e- → share or gain 1 e- and they tend to form 1- ions.

15. Noble Gases Last naturally occurring elements to be discovered b/c they are colorless & unreactive. Very stable with full valence electrons = 8. –(except He w/ 2) With lots of energy you can get Xe, Kr and Ar compounds. (No known He or Ne compounds) –In 1962 the first compound of the noble gases was prepared: XeF 2, XeF 4, and XeF 6. (flexible e- arrangments b/c f orbitals) –To date the only other noble gas compounds known are KrF 2 and HArF.

16. Transition elements (metals) d-block f-block u Inner transition elements