1. Examination of properties reveals why
Why “periodic?”
Examination of properties reveals why

2. Learning objectives Define ionization energy and electron affinity
Describe periodic trend in atomic and ionic radius and ionization energy
Predict order of atomic/ionic sizes using concept of shielding and periodic table

3. Properties show periodic variation

4. Periodic trends in atomic size

5. Atoms and ions Ions are created by removing or adding electrons
Positive ions are smaller than the neutral atoms
Negative ions are larger than the neutral atoms

6. Isoelectronic ions Isoelectronic ions have same number of electrons
Na [Ne]3s1; Mg [Ne]3s2; Al [Ne]3s23p1
Na+ [Ne]+; Mg2+ [Ne]2+; Al3+ [Ne]3+
P [Ne]3s23p3; S [Ne]3s23p4; Cl [Ne]3s23p5
P3- [Ar]3-; S2- [Ar]2-; Cl- [Ar]-
Isoelectronic cations, higher charged ions are smaller (nuclear attraction is stronger)
Na+ > Mg2+ > Al3+
Isoelectronic anions, higher charged ions are larger (nuclear attraction is weaker)
P3- > S2- > Cl-

7. Ionization energy: energy required to remove electron from isolated gaseous atom: A(g) = A+(g) + e

8. Electron affinity: energy released when electron is added to isolated gaseous atom: A(g) + e = A-(g)

9. Explain these trends
Atomic radius decreases across period, even though atomic number increases
Ionization energy increases – electrons more tightly held

10. Shielding and effective nuclear charge
The “shell” picture helps to explain these observations
Electrons in same shell experience stronger attraction to nucleus as shell fills
Nearly full – high charge
Nearly empty – low charge