1. Chapters 6 & 7 Chemistry 1L Cypress Creek High School
Unit 7: Periodic Table
Chapters 6 & 7
Chemistry 1L
Cypress Creek High School 1

2. Part 4: Periodic Trends 2

3. Periodic Table Trends Patterns on the periodic table Atomic Radius
Ionic Radius
Electronegativity
Ionization Energy

4. Periodic Trends Depend upon 4 important factors…
Energy levels – the horizontal rows; ranked from 1-7 based on energy and distance from the nucleus
Valence electrons – number of electrons in outermost energy level
Shielding effect – the decrease in the attraction between the outer electrons and the nucleus due to the presence of other electrons between them
Nuclear charge – depends on the number of protons – the more protons in the nucleus, the greater pull they have on their surrounding electrons

5. Atomic Radius
Atomic radius is half the distance between the centers of two atoms that are just touching each other
Influenced by 2 factors
The number of energy levels
The nuclear charge (pull of the positively charged nucleus on its electrons)
The more energy levels, the ________ the atomic radius.
(larger/smaller)
The more protons in the nucleus, the ________ the atomic radius.
larger
smaller

6. Atomic Radius Trend Atomic radius increases as you move down a group
Atomic radius decreases as you move from left to right in a period

7. Ionic Radius Metals lose electrons to form cationsLi
Li+
Atomic radius decreases - energy level is lost or “shed” (think of peeling an onion)
Nonmetals gain electrons to form anions
Make ions visibly smaller F F-
Atomic radius increases - energy level expands because it is more “crowded” and electrons exert greater repulsive forces on each other (think of 7 people vs. 8 people holding hands in a circle)

8. Ionic Radius Trend Ionic radius increases as you move down a group
Ionic radius decreases as you from left to right in a period BUT…
Energy levels change between cations and anions
Note: metals make smaller ions, nonmetals make larger ions
decreasing ionic radius

9. Atomic & Ionic Radius Trends9

10. Electronegativity
Electronegativity is a measure of how easily an atom attracts the valence electrons of another atom
Numbers are assigned to each element to rate the electronegativity (from 0.7 to 4.0)
Low electronegativity = does not want to attract valence electrons (metals)
High electronegativity = really wants to attract valence electrons (nonmetals)
Influenced by 2 factors:
Valence electrons
Shielding 10

11. Electronegativity Trend
Electronegativity decreases as you move down a group
Electronegativity increases as you from left to right in a period 11

12. Ionization Energy
Ionization Energy – the energy needed to remove the outermost electron in an atom
Influenced by 2 factors:
Nuclear charge – more protons pulling on the electrons, making it harder to remove them
Shielding – Radius is larger; outer electrons are farther from the nucleus; more difficult to gain electrons

13. Ionization Energy
First ionization energy is that energy required to remove the first electron
Ex: Easiest to remove Na’s first electron, hardest to remove Ar’s first electron
Second ionization energy is that energy required to remove the second electron
Ex: Easiest to remove Mg’s second electron, hardest to remove Na’s second electron
Third ionization energy is that energy required to remove the third electron
Ex: Easiest to remove Al’s third electron, hardest to remove Mg’s third electron
Fourth, fifth, sixth etc… the ionization energy patterns continues
This graph shows first ionization energy only!

14. Ionization Energy Trends
Ionization energy decreases as you move down a group
Ionization energy increases as you from left to right in a period
Increasing ionization energy
Decreasing ionization energy

15. Summary of Periodic Trends