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. Trend #2: Electronegativity
What is it?
The attraction that an atom has for a shared pair of electrons in a covalent bond.
It’s a measure of how good an atom’s nucleus is at the “tug-of-war” for shared electrons.
Most common scale used for electronegativity is the Pauling scale.
In the Pauling scale, electronegativity values range from

13. Why is there no value for He, Ne, Ar, and Kr?

14. Pauling Electronegativity

15. Examples
Explain why magnesium has a higher electronegativity than strontium.
Explain why chlorine has a higher electronegativity than silicon.

16. Explanation Across the period, electronegativity increases.
All elements in a period have their valence electrons in the same energy level.
The major difference is the increased nuclear charge as more protons are added to the nucleus.
Down the group, electronegativty decreases.
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.

17. Trend #3: Ionization Energy
What is it?
The amount of energy needed to remove an electron from the outermost shell.
The higher the ionization energy, the harder it is to remove the electron.
The “first ionization energy” refers to removing an electron from a neutral, gaseous atom.
Ex. F (g)  F+ + e-

19. First Ionization Energy Plot

20. Explanation Across the period, 1st ionization energy increases.
All elements in a period have their valence electrons in the same energy level.
The major difference is the increased nuclear charge as more protons are added to the nucleus.
Down the group, 1st ionization energy decreases.
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.

21. Related Concept: Subsequent ionization energies
The 2nd ionization energy is the energy required to remove an electron from a +1 ion.
The 2nd ionization energy for an element is ALWAYS larger than the 1st ionization energy. (And the 3rd I.E. is even larger, etc.)
Because there are now more protons than electrons, the protons have a stronger pull on the valence electrons.

22. Subsequent ionization energies
An element that has two valence electrons will have:
A relatively small 1st I.E.
A slightly larger 2nd I.E.
A MUCH, MUCH larger 3rd I.E.
WHY?

23. First, Second, Third, and Fourth Ionization Energies of Sodium, Magnesium, and Aluminum (kJ/mol)
1st IE
2nd IE
3rd IE
4th IE Na
496
4562
6912
9543 Mg
738
1451
7733
10540 Al
578
1817
2745
11570