1. Periodic Relationships Among the Elements

2. General Periodic Trends Atomic and ionic sizeAtomic and ionic size Ionization energyIonization energy ElectronegativityElectronegativity Electron AffinityElectron Affinity Higher effective nuclear charge Electrons held more tightly Larger orbitals. Electrons held less tightly. Shielding Effect!

3. Effective nuclear charge (Z eff ) is the “positive charge” felt by an electron. Na Mg Al Si 11 12 13 14 10 1 2 3 4 186 160 143 132 Z eff Core Z Radius Z eff  Z – number of inner or core electrons 8.3

4. Effective Nuclear Charge (Z eff ) 8.3 increasing Z eff

5. 8.3

7. Atomic Radii 8.3

8. Atomic Size Size goes UP on going down a group.Size goes UP on going down a group. Because electrons are added further from the nucleus, there is less attraction. This is due to additional energy levels and the shielding effect. Each additional energy level “shields” the electrons from being pulled in toward the nucleus.Because electrons are added further from the nucleus, there is less attraction. This is due to additional energy levels and the shielding effect. Each additional energy level “shields” the electrons from being pulled in toward the nucleus. Size goes UP on going down a group.Size goes UP on going down a group. Because electrons are added further from the nucleus, there is less attraction. This is due to additional energy levels and the shielding effect. Each additional energy level “shields” the electrons from being pulled in toward the nucleus.Because electrons are added further from the nucleus, there is less attraction. This is due to additional energy levels and the shielding effect. Each additional energy level “shields” the electrons from being pulled in toward the nucleus.

9. Atomic Size Size decreases across a period owing to increase in the effective nuclear charge. Large Small

10. Cation is always smaller than atom from which it is formed. Anion is always larger than atom from which it is formed. 8.3

13. Ionization energy is the minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state. I 1 + X (g) X + (g) + e - I 2 + X + (g) X 2 + (g) + e - I 3 + X 2+ (g) X 3 + (g) + e - I 1 first ionization energy I 2 second ionization energy I 3 third ionization energy 8.4 I 1 < I 2 < I 3

14. 8.4

15. Filled n=1 shell Filled n=2 shell Filled n=3 shell Filled n=4 shell Filled n=5 shell 8.4

16. General Trend in First Ionization Energies 8.4 Increasing First Ionization Energy

17. Electron affinity is the negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion. X (g) + e - X - (g) 8.5 F (g) + e - F - (g) O (g) + e - O - (g)  H = -328 kJ/mol EA = +328 kJ/mol  H = -141 kJ/mol EA = +141 kJ/mol

18. 8.5

20. Group 1A Elements (ns 1, n  2) M M +1 + 1e - 2M (s) + 2H 2 O (l) 2MOH (aq) + H 2(g) 4M (s) + O 2(g) 2M 2 O (s) Increasing reactivity 8.6

21. Group 7A Elements (ns 2 np 5, n  2) X + 1e - X - 1 X 2(g) + H 2(g) 2HX (g) Increasing reactivity 8.6

22. Group 8A Elements (ns 2 np 6, n  2) 8.6 Completely filled ns and np subshells. Highest ionization energy of all elements. No tendency to accept extra electrons.

23. Properties of Oxides Across a Period basicacidic 8.6 More basic Less basic More acidicLess acidic

24. Electronegativity,   is a measure of the ability of an atom in a molecule to attract electrons to itself. Concept proposed by Linus Pauling 1901-1994 Concept proposed by Linus Pauling 1901-1994

25. Periodic Trends: Electronegativity In a group: Atoms with fewer energy levels can attract electrons better (less shielding). So, electronegativity increases UP a group of elements. In a period: More protons, while the energy levels are the same, means atoms can better attract electrons. So, electronegativity increases RIGHT in a period of elements.

26. Electronegativity