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

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

Properties show periodic variation

Periodic trends in atomic size

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

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-

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

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

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

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