1. Periodic Trends OBJECTIVES:
Interpret group trends in atomic radii, ionic radii, ionization energies, and electronegativities.
Interpret period trends in atomic radii, ionic radii, ionization energies, and electronegativities.

2. Trends in Atomic Size
Where do you start measuring from?

3. Atomic Size }
Radius
Atomic Radius = half the distance between two nuclei of atoms
in the solid state (by X-ray diffraction)
or of a diatomic molecule.

4. Atomic Size Influenced by three factors: Energy Level
Higher energy level is further away.
Charge on nucleus
More positive charge pulls electrons in closer.
Shielding effect
The inner electrons shield the outer electrons from the nuclear charge/attraction.

5. Shielding
The electron on the outermost energy level has to look through all the other energy levels to see the nucleus.
Increases down a Group, & is Constant across a Period.
Shielding Across a Group is Constant, but the EFFECTIVE NUCLEAR CHARGE INCREASES.

6. Group Trends in Atomic Size
As we go down a group… each atom has another
energy level, so the atoms get bigger.
Shielding increases as well, so the nucleus has less of a hold on e-… distance is longer.
The Increased size of the Energy Levels down a group outweighs the increased nuclear charge H Li Na K Rb

7. Periodic Trends in Atomic Size
As you go across a period, the radius gets smaller.
More nuclear charge.
Outermost electrons are closer. Na Mg Al Si P S Cl Ar

8. Trends in Ionization Energy
The amount of energy required to completely remove an electron from a gaseous atom.
The energy required to remove the first electron is called the first ionization energy.

9. Ionization Energy
The second IE is the energy required to remove the second electron.
Always greater than first IE.
The third IE is the energy required to remove a third electron → Greater than 1st or 2nd IE.

10. What Affects the IE
The greater the nuclear charge, the greater the IE.
Larger positive nucleus has a greater attraction for the electrons, so the IE increases.
Greater distance from nucleus decreases IE
Electrons are further away from the attractive nucleus, and are easier to remove.

11. Group trends on IE As you go down a group, first IE decreases because…H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87
As you go down a group, first IE decreases because…
The electron is further away.
More shielding.

12. Periodic trends on IE
Across the representative elements:
The atoms are in the same period & have the same energy level.
Same shielding.
But, increasing nuclear charge holds e-’s tighter.
So IE generally increases from left to right.****
Full Energy Levels require lots of energy to remove their electrons.
Noble Gases have full orbitals.
Atoms behave in ways to achieve noble gas configuration.

13. Trends in Electron Affinity
What is Electron Affinity?
It’s the energy change associated with adding an electron to a gaseous atom.

14. Trends in Electron Affinity
It’s easiest to add an electron to Group 7A.
It gets them to a full energy level, or completes the OCTET.
Increase from left to right: atoms become smaller, with greater nuclear charge.
Decrease as we go down a group.

15. Electron Affinity in the Periodic Table

16. Group Trends Li1+ Na1+ K1+ Rb1+ Cs1+
Going down a Group, you are adding energy levels
Ions get bigger as you go down.
Li1+
Na1+
K1+
Rb1+
Cs1+

17. Electronegativity
The tendency for an atom to attract electrons to itself when it is chemically combined with another element.

18. Electronegativity Group Trend
The further down a group, the farther the electron is away, and the more electrons an atom has.
Pull/Attraction of the positive nucleus is lessened due to increased distance and Shielding.
Electronegativity decreases.
More willing to share.

19. Electronegativity Period Trend
As you move across a Period, there are the same number of energy levels, the same shielding, however …
Pull/Attraction of the positive nucleus on other’s electrons increases as the nucleus gets larger
Electronegativity Increases.

21. Periodic Properties Lab & Lab Write-Up