Lecture 7

1. Lecture 7 Ions and Ionic Compounds 2.1, C.1-C.3 and D.1-D.3 6-September Assigned HW C.7, C.10, C.12c,d, D.2a,b, D.4a,b, D.15, D.16, 2.8, 2.16, 2.18, 2.22, 2.24, 2.26, 2.28 Due: Monday 13-Sept

2. Review 1.11-1.18 • When an atom has multiple electrons, there are multiple components to the net coulombic potential that describes any one electron. • Shielding of the (+) charge of the nucleus occurs • We call the charge that an electron actually feels effective nuclear charge, Zeff. • Shielding also influences the energy of the orbitals within the same shell • s < p < d < f • Hund’s Rules tells us how to fill atomic orbitals • Add Z electrons with the lowest energy orbital occupied 1st. 2 electrons per orbital maximum • Fill degenerate orbitals such that each orbital has one electron with parallel spin before adding a 2nd electron to any orbital. • We can use shorthand notation to simplify our lives. • Atomic radius increases with shell and decreases across a period • Ionic radius can be predicted when considering the number of electrons vs. nuclear charge • Ionization energy scales inversely with atomic radius (small radius, big ionization energy.

3. Compounds Compound Bond Type Inorganic Compound Organic Compound Covalent Composed of Charged components Ionic Bonds Contains Carbon Cations Anions

4. Periodic Trends – Atomic Radius • Atomic Radius is dictated by the number of shells and the charge of the nucleus vs. total charge of electrons. • Shell – The atomic radius increases with the shell – so increases and we go DOWN the periodic table • Example: The atomic radius of K > Na • Within the same shell, we’re adding protons and electrons • Shielding is still ~the same, so increasing Z decreases radius • Example: The atomic radius of Cl < S < P General Trend: decreases from left to right and increases down

5. Periodic Trends – Atomic Radius Decreasing radius

6. Periodic Trends – Ionic Radius • What influence will adding an electron have to the radius of an atom? • Have we added any protons? • So we have NO additional attraction  radius increases • What happens if we remove an electron? • Remove repulsive force from electron-electron interactions • Remaining electrons feel more (+)  radius decreases

7. Periodic Trends – Ionic Radius Same # electrons, so why different size?

8. Ionization Energy and Electron Affinity • Ionization Energy: • The amount of energy to remove an electron from an atom • We’ve seen this concept before as work function in the photoelectric effect • 1st ionization energy < 2nd ionization energy, etc. • Electron Affinity: • This value describes how much an atom want to gain an electron. • These values are inversely related to atomic radius • When the radius is small, an electron is bound tightly, so high ionization energy. • Similarly, small radius = Zeff, so easier for another electron to be added.

9. Forming Ions - Cations • In general, when atoms form ions, they try to gain a full outer shell • s-block is typically found as cations – and REALLY want to be cationic. What charge would Mg likely form? Mg 3s [Ne] Oxidation State  charge on an ion

10. Forming Ions - Cations How about Al? Al 3p 3s [Ne]

11. Ionization Energies and Cations Do these trends make sense?

12. Forming Ions - Anions • Anions want that cozy full shell feeling too. Which block is most likely to form anions? What charge would a Cl ion likely have? Cl 3p 3s [Ne]

13. Forming Ions - Anions How about P? Which of these ions would have a larger ionic radius? P 3p 3s [Ne]

14. Forming Ions – Transition Metals • Predict a charge for Zn (Z = 30) • Now Fe (Z = 26)

15. Ionization Energies Revisited 5s 4p 3d 4s [Ar] There is a trend observed in the middle and of each block

16. Common Ions of Transition Metals • This special stabilization that occurs when degenerate orbitals are half full and completely full, some atoms are present with charges we would not necessarily predict. Zn  [Ar]4s23d10 Zn2+ [Ar] 3d10 3d 3d 4s 4s 2e- [Ar] [Ar]

17. Common Ions of Transition Metals - Fe Fe  [Ar]4s23d6 Fe2+ [Ar] 3d6 3d 3d 4s 4s 2e- [Ar] [Ar] Fe3+ [Ar] 3d5 e- 3d 4s [Ar]

18. Common Oxidation States - Carbon C  [He] 2s2 2p2 C2+ [He] 2s2 C4+ [He] 2p 2p 2p 2e- 2e- 2s 2s 2s [He] [He] [He] e- 3e- C4- [He] 2s2 2p6 C- [He] 2s2 2p3 Half full 2p 2p 2p 2s 2s [He] [He]

19. Common Oxidation States - Carbon • Multiple stable oxidation states for carbon. • +2  [He] 2s2 • +4 [He] • -1 [He] 2s2 2p3 • -4 [He] 2s2 2p6 • Which will be the most stable? Movie on common oxidation states of elements

20. Naming Compounds – Common Ions You can predict the charges on these Learn these!

21. Naming Compounds – Common Ions Learn these!

22. Ionic Compounds • Multiple ions can interact to form ionic compounds • Just like energy or matter, charge must be conserved – • Electrostatic Neutrality If Mg2+ and P3- interact, what do they form? • Find the lowest common multiple  2*3 = 6 • We need a total charge of 6 from each ion to make a neutral compound Na+ + Cl- NaCl +2 -3 +2 -3 +2

23. Naming Ionic Compounds • Naming rules: Example: CrCl3 • Identify the cation and anion and determine the charge • Name the cation. If the cation is a transition metal, identify the charge by converting it to a Roman numeral and placing it in parentheses. • Name the anion. • Combine the names – cation goes first.

24. Naming Ionic Compounds Mn2O3 Identify the cation and anion and determine the charge Name the cation. If the cation is a transition metal, identify the charge by converting it to a Roman numeral and placing it in parentheses. Name the anion. Combine the names – cation goes first. Cation Anion  Why don’t we indicate the number of cations or anions?

25. Naming Ionic Compounds - Hydrates Mn2O3 • 7H2O • When the ionic compound is hydrated (H2O is in the formula), it is named by adding hydrate with the appropriate greek prefix after the name.