Why “periodic?”

1. 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