1. Periodicity, Bond types
AP Chemistry
Mrs. Crowley

2. Father of Periodic table
Correctly predicted elements that had not been discovered yet
Dmitri Mendeleev
Based on “relative mass” comparison with Hydrogen (lightest element).

3. How’d he do that? Periodic law:
Properties of elements on the periodic table REPEAT themselves. Sometimes called “Mendeleev’s law.”
The chemical and physical properties of the elements in the same family are similar.
3 Major trends you need to know are….

4. Bond type depends on how electrons are shared between atoms
Electronegativity: An atom’s attraction for electrons
High EN = strong attraction
Low EN = low attraction

5. 3 types of bonds Ionic: very strong, creates solids.
Polar: medium strength bond. Some solids, liquids, and gases
Covalent: very weak bond. Some solids, liquids, gases

6. How to figure out bond type:
You must calculate at the electronegativity difference between the two atoms
Use the approximate guideline for electronegativity differences:
0.0 – < 0.5 = covalent (electrons shared nicely)
0.5 – < 2.0 = polar covalent (electrons being pulled from stronger atom)
2.0 – 4.0 = ionic (electrons stolen/transferred by stronger atom

7. Examples: H and Cl 2.1 – 3.0 Difference = 0.9 = polar C and H
2.5 – 2.1
Difference = 0.4 = covalent

8. Examples: Na and Cl 0.9 – 3.0 Difference = 2.1 = ionic
Also notice that you have a metal bonded to a nonmetal!
Follow this link for a video recap:
Y7iU

9. Use these values provided to help:
EN increases
EN decreases

10. Atomic Size

11. Complicated area to explain: Coulomb’s and Aufbau
Atomic Radius
Complicated area to explain: Coulomb’s and Aufbau
4th
1st
2nd
3rd
*Note: 100 pm = 1Å

12. Why?
As you go across a period, the principal energy level (# of rings) remains the same. As the nuclear charge (# protons) increases, it draws those electrons in closer to the nucleus. Atomic size therefore decreases going across the periodic table,
Atomic size increases when you move down the periodic table because more rings of electrons are being added.

13. Ionization Energy
Energy required to remove an electron from a gaseous atom. This causes a +1 charge.
The energy required to remove the first outermost electron is called the 1st ionization energy.
The larger the atom is, the easier it is to rip away an electron. Electrons that are far away from the nucleus are easier to pull away.

14. Ionization Energy Trends
Increases
Decreases

15. Trends in 1st ionization energy