1. The Periodic Table

2. Dimitri Mendeleev dev. the modern periodic system
Placed elements in order of increasing atomic mass
Reversed the order of some elements to keep similar elements in the same column
Ex. Ar and K
Predicted the existence and properties of elements that had not yet been discovered-this demonstrated the value of the table

3. Modern Periodic Table
Elements are placed in order of increasing atomic number
Periodic Law:
- the phys. and chem. properties of the elements are periodic functions of their atomic numbers.

4. Overview of Periodic Table
Horizontal rows are called periods
Elements in a period have similar electron configurations.
Ex.
Li -1s22s1
B- 1s22s22p1
**the 2nd energy level is the outermost for both
Vertical columns are called groups-elements in a group have similar properties
Ex. Group 1 elements are all highly reactive

5. The table is divided into two distinct regions: metals and nonmetals
Metals are in green, nonmetals in yellow, and
metalloids in blue.

6. Properties of Metals Good conductors of heat and electricity
Most are solids at room temp.- the only exception is Hg
Malleable-can be bent
Ductile- can be drawn into a wire

7. Properties of Nonmetals
Poor conductors of heat and electricity
Not malleable or ductile
Most are gases or liquids at room temp.

8. Properties of Metalloids
Have properties of both metals and nonmetals
Often used as semiconductors

9. Main Group Elements Groups 1, 2, and 13-18
Sometimes called the representative elements.

10. Transition Elements
Located in groups 3-12

11. Alkali Metals
React w/water to produce alkaline solutions and have metallic properties
Soft
Have shiny surfaces
Highly reactive
Become plasmas in gaseous state
All located in Group 1

12. Alkaline Earth Metals Harder, denser, and stronger than alkali metals
Less reactive than alkali metals
All located in Group 2

13. Halogens
Halogen is derived from a Greek word that means “salt former”
These elements combine w/metals to form salts
Most reactive group of nonmetals
All located in Group 17
Chlorine
Bromine

14. Noble Gases
Formerly known as inert gases because they were believed to be completely unreactive
Are unreactive in nature because of their full outermost energy levels
All located in Group 18

15. Lanthanides and Actinides
Collectively known as the Rare Earth Metals
Lanthanides have atomic #’s are shiny, reactive metals
Actinides have atomic #’s and are all radioactive

16. Hydrogen The most common element in the universe
Unique behavior due to the fact that it only has 1 proton and 1 electron
Can react w/many elements

17. Periodic Trends

18. Many properties of elements change in predictable ways:
There are six periodic trends we will talk about.
Atomic Radii
Ionic Radii
Ionization Energy
Electron Affinity
Electronegativity
Valence Electrons

19. Atomic Radius
The distance from the center of an atom’s nucleus to its outermost electron
Atoms get larger down a group
Atoms get smaller across a period
Why?
As you move down a group, there are more principle energy levels.
As you move across a period, the attraction between the increasing # of protons and electrons causes the atom to shrink.

20. Note: You have to ignore the noble gas at the end of each period.
Because neon and argon don't form bonds,
you can only measure their van der Waals radius
- a case where the atom is pretty well "unsquashed"

21. Examples Who has the larger radius: Li or Rb Ca or Br Be or Ba
Al or Ga
Answers:
1. Rb
2. Ca
3. Ba
4. Ga

22. Ionic Radii One-half the diameter of an ion in an ionic compound
Cation- positive ion
Anion-negative ion
The formation of a cation results in a loss of one or more electrons and leads to a decrease in radius.
The formation of an anion results in a gain of one or more electrons and an increase in radius.

23. Examples: Which has the larger radius? Ca or Ca+2 Br or Br-1
Fe+2 or Fe+3
Answers: Ca
Br-1
Fe+2

24. Ionization Energy The energy required to REMOVE AN ELECTRON.
Why??
The increased pull of additional protons as you go across a period results in a stronger pull on electrons which means more energy to remove them.
Down a group, the outer electrons are further from the pull of the protons in the nucleus which means less energy to remove them.
The energy required to REMOVE AN ELECTRON.
Increases across a period
Decreases down a group

25. The ionization energy for the removal for a second or third electron will be higher.
Examples:
Who has the larger ionization energy?
Ca or Kr
Be or Ba
Fe+2 or Fe+3
Answers: Kr Be
Fe+3

26. Electron Affinity
The energy change that occurs when an atom gains an extra electron.
Becomes more negative across a period ( the more negative, the more likely an electron will be gained)
Becomes less negative down a group ( meaning less likely to gain an electron).

27. Examples:
Answers: Cl F
Which has the more negative electron affinity and therefore is more likely to gain an electron?
Mg or Cl
F or At

28. Electronegativity Ability to attract electrons.
Fluorine is the most electronegative element.
Increases as you go across a period.
Decreases down a group.
Which has the higher electronegativity meaning it is more likely to gain an electron?
F or Ne
Cl or I
Answers:
F ( the trend stops at Noble Gases) Cl

29. H: 2. 1 Li: 1. 0 Be: 1. 5 B: 2. 0 C: 2. 5 N: 3. 0 O:3. 5 F: 4. 0 Na: 0
H: 2.1 Li: Be: 1.5  B: 2.0 C: 2.5 N: O:3.5 F: Na: 0.9 Mg: 1.2 Al: 1.5 S: 1.8 P: 2.1 S: Cl: 3.0

30. Valence Electrons Group # Valence 1 1 2 2 13 3 14 4
1 1
2 2
13 3
15 5
16 6
17 7
18 8
The electrons available to be lost, gained, or shared in the formation of a chemical bond-the outermost electrons.