1. The Periodic Table
Chapter 6

2. Periodic Table
Arrangement of elements in which the elements are separated into groups based on a set of repeating properties
Allows you to easily compare the properties of one element (or groups of elements) to another element (or groups of elements)
Period: horizontal row
Group (family): vertical column
Number and letter (Group 1A)

3. Organizing the elements
Mendeleev published his version of periodic table in 1869
He arranged the elements in order of increasing atomic mass.
Figure 6.3 page 156
He left spaces predicting that elements would be discovered to fill in the holes and he predicted properties based on location on periodic table.
Now elements are arranged by increasing atomic number

4. Periodic law
The pattern of properties within a period repeats as you move from one period to the next

5. Metals Good conductors of heat/electricity High luster/sheen
Reflects light
Solid at room temp (except Hg)
Ductile (pulled into wires)
Malleable (hammered into thin sheets)
Alkali metals: Group 1A (soft, explode in water)
Alkaline earth metals: Group 2A (very reactive)
Transition metals: B groups

6. Nonmetals Upper right hand corner Most are gases at room temperature
Poor conductors heat/electricity (C is exception)
Brittle (shattered when hit with hammer)
Halogens (salt-formers): Group 7A (all 3 states of matter)
Noble gases: Group 8A (inert gases)

7. Metalloids Border the stair-step
Properties similar to metals and nonmetals
Boron, Silicon, Germanium, Arsenic, Antimony, Tellurium, Astatine

8. Coloring the Periodic Table
Alkali metals: red
Alkaline earth metals: dark orange
Transition metals: light orange
Inner transition metals: yellow
Other metals: light pink
Halogens: dark purple
Noble gases: blue
Other nonmetals: light purple
Outline stair step: black
Metalloids: green
Key of colors

9. Metallic characteristics
Metals want to lose electrons, not gain
As you move from left to right, atoms tend to gain electrons, therefore, decreasing the metallic characteristics.
Metals are to the left of the staircase and nonmetals are to the right of the staircase.

10. Atomic Size
½ the distance between the nuclei of two atoms of the same element when the atoms are joined
Nuclear charge: strength of the nucleus’ pull
Based on the number of protons in the nucleus
Carbon as a nuclear charge of 6+
Shielding effect: electrons between outer electrons and nucleus shield the outer electrons from the full nuclear charge of the nucleus

11. Group trends in atomic size
As atomic # increases within group, the charge on the nucleus increases and number of occupied energy levels increases.
Increase in positive charge draws electrons closer to nucleus.
Shielding effect is greater than the effect of the increase in nuclear charge
Atomic size increases

12. Periodic trends in atomic size
Across a period, electrons are added to the same principal energy level
Shielding effect is constant for all the elements in the period
Increasing nuclear charge pulls electrons in highest energy levels closer to nucleus
Atomic size decreases

13. Ions
Positive and negative ions form when electrons are transferred between atoms
Cation: ion with positive charge
Metals
Loses electrons
Na1+
Atom is larger than cation
Anion: ion with negative charge
Nonmetals
Gains electrons
Cl1-
Atom is smaller than anion

14. Ionic Radius Radius of the ion.
Cations are smaller than the atom they form from.
Anions are larger than the atom they form from.

15. Ionization Energy Energy required to remove an electron from an atom
Gaseous state
First ionization energy is energy required to remove 1st electron
Second ionization energy is energy required to remove the second electron

16. Electronegativity
Ability of an atom to attract electrons when the atom is in a compound
Noble gases do not have electronegativity because they do not form compounds
Table 6.2 page 177

17. Electron Affinity
Change in energy of a neutral atom in a gaseous state when an electron is added to form a negative ion
Likelihood of an atom gaining an electron