1. Module 3.03
Periodic Trends

2. Periods – rows of the periodic table
To Review:
Periods – rows of the periodic table
Groups – columns of the period table

3. Periodic Table

4. Trends are useful in making predictions
Periodic table is arranged based on the properties of the elements
Reading across each period and down each group you will see repeated trends in properties

5. Effective Nuclear Charge
Effective Nuclear Charge – the charge (from the nucleus) felt by the valence electrons after the number of shielding electrons are taken into account. Subtract the number of shielding (core) electrons from the total nuclear charge (number of protons) gives the effective nuclear charge.

6. Effective Nuclear Charge
Example:
Nitrogen
7 protons (atomic number of 7)
-2 core electrons (1s2 electrons)
5 is the effective nuclear charge
Each of the valence electrons in nitrogen feels an effective nuclear charge of about +5.

7. Effective Nuclear Charge
Across a Period:
The effective nuclear charge felt by an atom’s valence electrons increases by one for each element from left to right in a period
Down a Group:
The effective nuclear charge felt by an atom’s valence electrons stays constant for each element going down a group

8. Atomic Radius
Atomic Radius – half the distance between the centers of two atoms that are bonded together.

9. Atomic Radius
Across a Period:
There is a gradual decrease in atomic radii from left to right across a period.
Down a Group:
There is a general increase in atomic radii going down each group on the periodic table.

10. Atomic Radius

11. Ionization Energy
Ionization energy – energy required to remove on electron from an element, resulting in a positive ion.
Elements with a lower effective nuclear charge felt by their electrons will give up an electron easier than other elements.

12. Ionization Energy
Across a Period:
Ionization energy has a general increase for elements across a period from left to right.
Down a Group:
Ionization energy of elements decreases going down a group because the atomic radius of the atoms increases.

13. Ionic radius – ½ the diameter of an ion.
Note: Metals naturally form cations (positive ions) by losing one or more electrons. Nonmetals naturally form anions (negative ions) by gaining one or more electrons.

14. Atomic Radius vs Ionic Radius

15. Ionic Radius
Across a Period:
Within each period, the metals at the left form cations and the nonmetals at the right form anions. There is a decrease in the ionic radii of the cations from left to right and a decrease in the ionic radii of the anions from left to right.
Down a Group:
Ionic radii increase down a group, following the same trend as atomic radii.

16. Electronegativity
Electronegativity – measure of the attraction of an atom for the electrons in a chemical bond.

17. Electronegativity
Across a Period:
Electronegativity increases from left to right across a period because of the effective nuclear charge.
Down a Group:
Electronegativity decreases down a group as a result of a greater atomic radius.

18. Module 3.03
Periodic Trends