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The periodic table - an arrangement of the elements in a tale according to their periodic properties.  The chemists most important tool in understanding and predicting chemical reactivity.  It groups elements in a manner that allows us to understand how they will react with other elements.

Go Have a look at the periodic table across from the chemistry office!  Also browse the web elements.

Father of the periodic table - Dimitri Mendeleev (1834-1907) - (brief excerpt from article read in class) - organized by reactivity and atomic mass.

Mendeleev created an early periodic table and even predicted the existence and properties of certain elements that had yet to be discovered.    For instance, here is his predictions for an element he called "Eka-silicon" and a comparison with the actual element now called Germanium

Property

Predicted for Eka-silicon (E)

Observed for Germanium

atomic mass (amu)

72

72.59

density (g mL-1)

5.5

5.32

appearance

dark gray

gray-white

oxide

EO2, white solid, density 4.7 g mL-1

GeO2, white solid, density 4.23 g mL-1

Navigating the periodic table (Show periodic table from you book)

 

1. Horizontal rows - periods
2. Vertical columns - groups

 


3. "A" Groups - main group or representative elements
4. "B" Groups - transition elements

 

Three main types of elements (Show Fig 4.2 from your book)

1. Metals - Left of solid zigzag line - generally solids, shiny, ductile (can be drawn into wires), malleable (can be beaten into thin sheets), high electrical and thermal conductivity

2. Nonmetals - right of solid zigzag line  - solids, liquids and gases, poor conductors, solids brittle

3. Metalloids - on the zigzag line - share properties of both metals and nonmetals - some technologically very important such as Si (semiconductor).

 

Modern Periodic Law - when  the elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.

 Note: Mendeleev used atomic mass, why does atomic number make more sense?

Periodic Table and Electron Configuration - Similar elements have similar electron configurations.

 

Let's look at electron configuration of group VIIIA (0 in the book) "Noble Gases"

2He   1s2

10Ne  1s22s22p6

18Ar  1s22s22p63s23p6

All have filled energy levels! - Such filled shell electron configurations are particularly stable, and consequently, the noble gases are among the least reactive elements.

Rule: for any given group in the periodic table, the electrons in the outermost energy shell ("valence" electrons) occupy the same types of orbitals. 

It is this similarity that gives them similar properties.

Group 1A - Alkali metals

 

One electron in the outermost shell (one valence electron)

3Li = 1s22s1

11Na = 1s22s22p63s1 or [Ne]3s1 for short and to emphasize the outer shell

19K = 1s22s22p63s23p64s1 or  [Ar]4s1

37Rb              [Kr]5s1

Group 2A - Alkaline earth metals

 

4Be  = [He]2s2

12Mg = [Ne]3s2

20Ca = [Ar]4s2 , etc...

Elements in groups (columns) of the periodic table share the same outer shell electron configuration (just a different primary energy level involved). (Fig 4.6)

Group 1A and 2A = s-block
Group 3A-7A = p-block
Group 6A - chalcogenides [ ]ns2np4
Group 7A - halogens [ ]ns2np5

"B" Groups - transition metals = d-block

Lanthanides and Actinides = f-block

A good measure of whether you understand the organization of the periodic table and electron configuration is to be able to draw the periodic table (just empty boxes).