1. Periodic Trends

2. Why study periodic trends?
There are numerous factors that affect the properties of elements and compounds.
We’ve looked at the forces between particles.
These forces are affected by factors such as the size of atoms, the size of ions, and ionization energy.

3. Explaining Periodic Trends
Periodic trends are related to how strongly the nucleus can attract valence electrons.
There are two major factors to consider:
The effective nuclear charge (Zeff) of the nucleus.
Zeff = # protons - # kernel electrons
The higher the Zeff, the stronger the attraction of the nucleus on the electrons.
The number of full levels of electrons between the nucleus and the valence electrons (shielding effect).
More shielding levels = weaker attraction of the nucleus on the electrons.

4. + Shielding Effect - - - - nucleus Electron
Valence + -
nucleus - -
Electrons -
Electron
Shield
“kernel”
electrons
Kernel electrons block the attractive force of the nucleus from the valence electrons

5. Trend #1: Atomic Radii

6. Relative Size of Atoms

7. Atomic Radius
Since the edge of an electron cloud is difficult to define, scientists use covalent radius, or half the distance between the nuclei of 2 bonded atoms.
Atomic radii are usually measured in picometers (pm) or angstroms (Å). An angstrom is 1 x m.

8. Covalent Radius
The nuclei of two Br atoms bonded together are 2.86 angstroms apart. So, the radius of each atom is 1.43 Å.
2.86 Å
1.43 Å

9. As you move across the period, the atomic radius decreases.
All elements in a period have their valence electrons in the same energy level. Therefore, the shielding effect is constant.
The major difference is the increased effective nuclear charge as more protons are added to the nucleus while the number of kernel electrons remains constant.
The attraction the nucleus has on valence electrons becomes stronger.
As a result, the atomic radius decreases.

10. As you move down the group, the atomic radius increases.
Atoms for elements in the same group have the same Zeff.
Major difference is the shielding effect.
More full energy levels between the nucleus and the valence electrons weaken the pull the nucleus has on those electrons.
As a result, the atomic radius increases.

11. What should your explanation look like?
Discuss both elements being compared.
State what makes them similar.
State what makes them different.
Explain how those statements answer the question.
Ex. Explain why atoms of oxygen are smaller than atoms of boron.
Ex. Explain why atoms of iodine are larger than atoms of bromine.

12. Related Idea - Ionic Radii
Our major concern when looking at ionic radii is being able to compare the radius of an ion to the radius of the atom from which it comes.

13. Metal atom vs. Metal ion (cation)

14. Metal atom vs. Metal ion (cation)
Cations are smaller than the atoms that they come from.
Always true: Same # of protons but fewer electrons. Result is a stronger pull from the nucleus.
Almost always true: All electrons are removed from outer energy level. (One example when this isn’t true is Sn2+)

15. Nonmetal atom vs. Nonmetal ion (anion)

16. Nonmetal atom vs. Nonmetal ion (anion)
Anions are larger than the atoms that they come from.
Always true: More electrons cause more forces of repulsion. This also causes the electrons to spread out and take up more space.
Always true: Same # of protons but more electrons. Result is a weaker pull from the nucleus.