1. Organization of the Periodic Table

2. Dmitri Mendeleev organized the elements in the first periodic table by order of increasing mass in 1870.
Found repetition in properties of elements
Widely accepted because it was able to predict the existence and properties of undiscovered elements
Medeleev’s table
Arranged by increasing atomic mass
Showed periodic pattern
Left blank spaces and made very accurate predictions about the missing elements and their properties

3. Problems with Mendeleev’s Table
Several elements were discovered after his table was published
The atomic masses for known elements were more accurately determined
It became apparent that several elements in his table were not in the correct order
Arranging the elements by mass put them in groups that did not have the same properties

4. In 1913 Henry Mosley, an English Chemist, reorganized the Periodic Table based on the number of protons.
Discovers that atoms of each element contain a unique number of protons in their nuclei
Number of protons = atomic number
Arranging elements in order of increasing atomic number solved the problems with elements in the table being out of order (as with atomic mass)
Showed a clear periodic pattern of properties
Led to the development of the periodic law

5. The Modern Periodic Table
Periodic law – there is a periodic repetition of chemical and physical properties of the elements when they are arranged by increasing atomic number

6. Elements are placed on the Periodic Table based on the number of protons in their nucleus

7. Arrangement of Table… families
Consists of boxes – each box represents a single element
Columns of boxes = groups (families)
Have the same number of valence electrons
Have similar end electron configurations
Have similar chemical and physical properties

8. Each element in a particular family contains the same number of outer shell electrons, or valence electrons.

9. Arrangement of Table cont’d … Periods
Rows of boxes = periods
Represent an energy level around the nucleus
period 1 = energy level 1
As you move across the period (left to right), there is an increasing number of valence electrons

10. Types of Elements Representative elements – Groups 1A-8A
Often referred to as the main group of elements
possess a wide range of chemical and physical properties
Transition elements – Groups 1B-8B
Also known as the transition metals
More recent numbering system – classifies groups as 1-18

11. Classifying Elements Other major divisions in the Periodic Table include the metals, nonmetals, and metalloids

12. Metals Nonmetals Metalloids
Silicon – used in solar cells and computer chips

13. Classifying Elements Metals Metalloids Nonmetals
Shiny, good conductors, solid at room temperature, ductile and malleable
Most elements on the periodic table are metals
Metalloids
Found on stair-step line separating metals and nonmetals in the periodic table
Only exception is Aluminum
Have chemical and physical properties of both metals and nonmetals
Nonmetals
Generally gases or brittle, dull-looking solids
Poor conductors of heat and electricity

14. Groups of Elements Alkali metals – Group 1A ( very reactive)
Alkaline earth metals – Group 2A (reactive)
Transition metals – B groups
Inner Transition metals – bottom two rows
Lanthanide and Actinide series
Boron Family – Group 3A
Carbon Family – Group 4A
Nitrogen Family – Group 5A
Oxygen Family – Group 6A
Halogens – Group 7A (extremely reactive)
Noble Gases – Group 8A (nonreactive)

15. Atomic Radius Atomic radius
How closely an atom lies to a neighboring atom
Trends within periods = decrease due to outermost e- being pulled toward nucleus
Trends within groups = increase due to shielding and increased energy levels

16. Atomic Radius

18. Ionic Radius Ionic radius
Ion=atom that has a + or – charge due to the gain or loss of e-
Lose e-/ + charge / smaller atom - electrostatic repulsion decreases & valence e- leaves unfilled orbital
Gain e-/ - charge / larger atom – electrostatic repulsion increases & causes increase distance between outer e- causing larger radius

20. Ionic radii
Trends within periods = size of the + ion decreases, then beginning in group 5A or 6A the larger – ion decreases
Trends within groups = ion size increases due to the ion’s outer e- being in higher energy levels

21. Ionization energy is the energy required to remove an electron from an atom

22. Ionization Energies

23. Ionization energy 1st, 2nd , 3rd etc.
How strong atom’s nucleus holds onto the valence e-
Large ionization energy, less likely to form + ions
Low ionization energy, atom loses outer e- easily and easily forms + ions

24. Ionization Energy Trends
Trends within periods = increase due to increase in nuclear charge producing an increased hold on valence e-
Trends within groups = decrease due to valence e- being farther from the nucleus requiring less energy to remove them

25. Octet Rule
Atoms tend to gain, lose, or share e- in order to acquire a full set of eight valence e-
Exception : 1st period completely filled with 2 e-
Elements on the right side of the periodic table tend to gain e- and tend to form – ions
Elements on the left side of the periodic table tend to lose e- and form + ions

26. Electronegativity is how much atoms pull electrons away from another atom

27. Electronegativity Values

28. Electronegativity Trends
Trends within periods = increase
Trends within groups = decreases
Lowest electronegativities are found at the lower left side of the periodic table
Highest electronegativities are found at the upper right side of the periodic table