1. Unit 6: The Periodic Table
Properties of Elements and Trends

2. Mendeleev The original periodic table was created by Dmitri Mendeleev
He organized the table by atomic mass
He predicted existence of 3 missing elements discovered later on

3. Modern Periodic Table
The modern periodic table is organized by atomic number.
Vertical columns  called groups or families organize elements by reactive properties (valence e-)
Horizontal rows  called periods also represent the shells or energy levels
The concept of periodic law is based on repeating patterns or trends in element properties.
Henry Mosely

4. Major Types of Elements
There are 3 main types of elements:
1) Metals 2) Metalloids 3) Non-Metals

5. Take Out Your Coloring Tools

6. Properties of Metals Solid at room temperature (except Hg)
Luster – Reflects light to give a shine
Malleable – Moldable or bendable
Ductile – Can be stretched into thin wire
Most have high melting points
Good conductor of heat & electricity
Have metallic bonds

7. Properties of Non-Metals
Can be solid, liquid or gas at room temp
Exist as Allotropes - different molecular forms of the same element (different properties, too) O2 and O3
Dull, lack luster
Brittle
Poor conductors

8. Ionic Compounds Metals Vs. Non-Metals
Metals lose electrons to form a full valence shell and become + ions (cations).
Non-metals gain electrons to form a full valence shell and become – ions (anions).
Ionic Compounds

9. Properties of Metalloids
Solid at room temperature
Semi-conductors
Have varying properties that can be metallic or non-metallic
6 of them:

10. Groups on the Periodic Table
In general, elements found within the same group have similar chemical and physical properties.
This is mainly due to the fact that they have the same number of valence electrons.

11. Alkali Metals (Group 1) Metals (very reactive)
Have one valence electron in outer shell
Tend to lose that one electron to become a stable cation.
Rarely found in nature.
Exist as solids at room temperature, but are soft
Silvery-white or grayish in color
Examples:

12. Alkaline Earth Metals (Group 2)
Metals (2nd most reactive)
Have two valence electrons
Tend to lose the two electrons to become a stable cation
Exist as solids at room temperature
Silvery white or grayish white in color
Examples:

13. Halogens (Group 17) Have 7 valence electrons
Non-metals (most reactive)
F, Cl (gases), Br (liquid), I (solid)
Tend to gain one electron to form a more stable anion.
Examples:

14. Noble Gases (Group 18)
Very stable because of filled valence shell (8 electrons).
All non-metal (gases at STP)
Not Reactive
Examples:

15. Transition Metals (Groups 3-12)
Also known as “heavy metals”
Not as reactive metals
Tend to be very dense (heavy for small amounts)
Have multiple oxidation states (can form more than one ion) so we use roman numerals when naming compounds with them.
Produce very colorful solutions when dissolved in water
Examples:

16. Lanthanide & Actinide Series
These are rare earth metals
Many have applications in nuclear technology
They actually fit in between groups 3 and 4 of the periodic table
Above 92 they are mostly man-made and are unstable

17. Periodic Trends
According to periodic law trends occur in the periodic table in the following Categories:
- reactivity
- metallic character
- atomic radius* & ionic radius*
- ionization energy*
-electronegativity*
*Sketch this on a page in your notebook

18. Period Trends in Atomic Radius
Atomic Radius – Is the distance from the center of the atom, “The Nucleus,” to the furthest away electron shell
Atomic radius decreases

19. Period Trends in Ionization Energy
The energy required to remove an electron from an atom is known as the FIRST IONIZATION ENERGY.
To remove a second electron requires what is called the SECOND IONIZATION ENERGY.
Increasing IE
Increasing IE

20. Cation Size
When an electron is lost from an atom, the nucleus now has more pull on the outer energy level making the ion SMALLER.
MINUS ONE ELECTRON

21. Anion Size
When an electron is gained the nucleus now has less pull on the outer energy level making the ion LARGER.
PLUS ONE ELECTRON

22. Electronegativity
Can be defined as an atom’s ability to attract electrons.
Expressed in arbitrary units on Table S of your Reference Tables.
Non-metals tend to have higher electronegativities because they like to gain electrons.
atoms
electrons