1. Chemical Periodicty
Patterns of Change

2. Topics of Discussion Periodic Table Electron Configurations
Periodic Trends

3. Development of the Periodic Table
Mid 1800’s
~ 70 elements known
- coinage metals
-examples:
-halogens: “salt formers”
-alkali metals
-examples
Early attempts at organizing the elements
- Döbereiner’s Triads
- Newland’s Law of Octaves

4. Döbereiner’s Triads
In 1829 Döbereiner proposed the Law of Triads: The middle element in the triad had atomic weight that was the average of the other two members. The densities of some triads followed a similar pattern. Soon other scientists found chemical relationships extended beyond triads. Fluorine was added to Cl/Br/I group; sulfur, oxygen, selenium and tellurium were grouped into a family; nitrogen, phosphorus, arsenic, antimony, and bismuth were classified as another group.

5. Dmitri Mendeleev Russian chemist
In 1869, published the first Periodic Table of the Elements
Dmitri Mendeleev

6. Mendeleev (cont)
Listed all known elements in order of increasing atomic mass
Arranged the columns that elements with similar properties were side by side…even if out of order by mass.

7. Mendeleev’s Predictions
Left blanks in the table if no element with appropriate properties was known
Accurately predicted properties of missing elements
“Ekaaluminum”
Gallium discovered in 1875
“Ekaboron”
-Scandium discovered in 1877
“Ekasilicon”
-Germanium discovered in 1886
Drawbacks: didn’t account for noble gases and rare earth elements

8. Henry Moseley Worked with Rutherford
Recorded X-ray diffraction data on the elements to determine nuclear charge

9. Moseley (cont)
Realized that the properties of the elements are related to the number of protons in nucleus
Rearranged the periodic table based on atomic number

10. Periodic Law
When the elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties
-regular and repeating pattern
metalmetalloidnonmetalunreactive gas

11. Valence Electrons The outermost electrons in an atom
lead to chemical properties
The electrons in the highest occupied s and p orbitals
varies from 1-8
number of valence electrons = group number
”main block”
groups 1, 2, 13-18

12. How many valence electrons?Li N O C Ar Mg

13. Lewis Dot Structures/Diagrams
One dot for each valence electron
Li C
N Ar
O Mg

14. Modern Periodic Table Periods in horizontal rows
numbered 1-7
Groups in vertical rows
numbered 1-18
aka; families
s block groups 1-2
p block groups 13-18
d block groups 3-12
f block lanthanides & actinides

15. s-, p-, d-, and f-blocks

16. Alternative Periodic Tables
Periodic table by Emil Zmaczynski (1935)

17. Alternative Periodic Tables
Spiral Periodic Table from Theodor Benfey
(1960)

18. Alternative Period Table
Timmothy Stowe’s Physicist Periodic Table (1988)

19. Atomic Size (Radius) Left right in a period size decrease
more protons, therefore electrons are held more tightly
Zpff increases left to right
Shielding constant in a period (number of core electrons is constant)

20. Atomic Radius As you go down a column, the radius increases in size
Outermost electrons are not held tightly by the nucleus

21. Ionic Radii Cations will be smaller than parent atom
Fewer electron electron repulsions
Anions will be larger than parent atom
More electron electron repulsions

22. Ionization Energy
From left to right, ionization energy increases in a period
Number of core electrons is constant
Number of protons increases
Down a group, ionization energy decreases
Increased shielding as you go down a column
Outermost electrons feel less nuclear charge and are easily removed

23. Electronegativity From left to right, electronegativity increases
F has the highest
Down a column, electronegativity decreases
Fr has the least

24. Summation Create a concept map from these three main concepts:
Periodic Table
Electron Configuration
Periodic Trends