1. Periodic Table Trends

2. Nuclear Charge (Effective Nuclear Charge)
Positive charge of nucleus increases as new electrons added are all at the same energy level.
Electrons feel a stronger pull towards the nucleus, and get closer (lower energy state)

3. Nuclear Charge (Zeff) (Effective Nuclear Charge)
(-) Electrons attracted to (+) nucleus
3 factors that effect this
The more protons in the nucleus, the greater the Zeff
The more distance between the nucleus and electrons  the smaller the Zeff
The more inner electrons that the atom have  the more repulsion will happen between the inner electrons and outer electrons (Valence Electron)  the smaller the Zeff

4. Atomic Radius
Atomic radius is half the distance between two bonded nuclei.
What is the trend as you move across a period?
Down a group?

5. Atomic Size/ Atomic radius
down groups  Increases
Electrons in increased energy levels (shells)  atomic radius is getting bigger
across periods  Decreases.
There are more valence electrons in the same shell (energy level) so they pull towards the nucleus which has an increased effective nuclear charge with more protons.  the distance will be closer  atomic radius is smaller.

6. Atomic Radius
Leaving the noble gases out, atoms get smaller as you go across a period.
They get larger as you move down a group.

7. Ionic Radius
The ionic radius of an element is its size after the atom has gain or lost electrons.

8. When atoms lose electrons  form (+) charge, atoms become smaller  loss of a valence electron  repulsion between electrons getting lower  allowing the electrons to be pulled closer to the nucleus.
When atoms gain electrons  form (-) charge, atoms become larger  get more electrons  repulsion between electrons getting increases and forcing them to move farther apart  increase distance between the outer electrons.

9. First Ionization Energy
Energy required to remove one electron from an neutral atom.

10. Ionization Energy (potential)
First Ionization Energy is the energy required to remove the most loosely held electron from one gaseous atoms.
X(g) + i.e. → X(g)+1 + e-
The second ionization energy is the energy required to remove the next most loosely held electron, etc…
What is the trend as you move across a period? Down a group?

11. Ionization Energy Down groups  Decreases Across periods  increases
as electrons are further from the nucleus less effective nuclear charge. electrons are removed more easily  less energy
Across periods  increases
Increased effective nuclear charge holds electrons more tightly. So if one electron wants to move  need more energy

12. Electronegativity-
Measure of ability of an atom to attract electrons from another atom in the compound

13. Trends in Electronegativity:
down groups  decreases
The electrons are further from the nucleus  less effective nuclear charge  less ability to attract electron from another atom.
across periods  increases
as atoms have more valence electrons in the same energy level and more protons in nucleus  they have a higher effective nuclear charge  easily to attract electron from another atom.

14. Electron Affinity
The energy change that occurs when an electron is added to the atom to form a negative ion.  gaining an electron.
A + e- → A- + energy
A + e- + energy → A-
Similar to electronegativity

15. Trends in Electron Affinity:
down groups  decreases
The electrons are further from the nucleus  less effective nuclear charge  less ability to attract electron  less gain an electron
across periods  increases
as atoms have more valence electrons in the same energy level and more protons in nucleus  they have a higher effective nuclear charge  easily to attract electron  easily to gain an electron.

16. Overall