1. CHEMISTRY 161 Chapter 8

2.  (n, l, m l, m s ) ATOMIC ORBITALS nlmlml orbitalsdesignation 10011s 20012s 1-1,0,+132p x,2p y,2p z 30013s 1-1,0,+133p x,3p y,3p z 2-2,-1,0,+1,+253d xy,3d yz,3d xz, 3d x 2 -y 2,3d z 2 4…………

3. H Atom Orbital Energies energy level diagram H atom 3s3p3d2s2p1s3s3p3d2s2p1s E energy depends only on principal quantum number orbitals with same n but different l are degenerate

4. 1s1s E 2s2s 2p2p 3s3s 3p3p 3d3d 4s4s 4p4p 5s5s 4d4d MULTI-ELECTRON ATOM orbitals with same n and different l are not degenerate energy depends on n and m l EXAMPLES [Xe] EXP I

5. Periodic Table of the Elements period groupgroup chemical reactivity - valence electrons ns 1 ns 2 ns 2 np 6 ns 2 (n-1)d x

6. PERIODIC TRENDS 3. IONIZATION ENERGIES 4. ELECTRON AFFINITIES 1. ATOMIC RADIUS 2. IONIC RADIUS

7. ATOMIC RADIUS MAIN GROUPS ATOMIC RADIUS EXP II

8. 1s, 2s, 3s 3s 2s 1s

9. 2p x, 3p x, 4p x 4p x 3p x 2p x

10. ATOMIC RADIUS MAIN GROUPS ATOMIC RADIUS effective nuclear charge (shielding s vs p orbitals)

11. IONIC RADII IONIC RADIUS

12. cations are smaller than their atoms anions are larger than their atoms Na is 186 pm and Na + is 95 pm F is 64 pm and F - is 133 pm same nuclear charge and repulsion among electrons increases radius one less electron electrons pulled in by nuclear charge O < O – < O 2–

14. EXAMPLES Which is bigger? Na or Rb Rb higher n, bigger orbitals K or Ca K poorer screening for Ca Ca or Ca 2+ Ca bigger than cation Br or Br - Br smaller than anion

15. QUESTION The species F -, Na +,Mg 2+ have relative sizes in the order 1F - < Na + <Mg 2+ 2F - > Na + >Mg 2+ 3Na + >Mg 2+ > F - 4Na + =Mg 2+ = F - 5 Mg 2+ > Na + >F -

16. QUESTION 1F - < Na + <Mg 2+ 2F - > Na + >Mg 2+ 3Na + >Mg 2+ > F - 4Na + =Mg 2+ = F - 5 Mg 2+ > Na + >F - Na + is 95 pm Mg 2+ is 66 pm F - is 133 pm ALL 1s 2 2s 2 2p 6 ALL are isoelectronic

17. 3. IONIZATION ENERGIES M(g)  M + (g) + e - energy required to remove an electron from a gas phase atom in its electronic ground state I 1 > 0 first ionization energy (photon)

18. M + (g)  M 2+ (g) + e - M 2+ (g)  M 3+ (g) + e - second ionization energy third ionization energy I 2 > 0 I 3 > 0 I 1 > I 2 > I 3

20. Why? electrons closer to nucleus more tightly held IONIZATION ENERGY first ionization energies decrease d shell insertion

21. IONIZATION ENERGY

23. 1. closed shells are energetically most stable 2. half-filled shells are energetically very stable DERIVATION OF IONIZATION ENERGIES

24. noble gases have the highest ionization energy

25. 4. ELECTRON AFFINITIES the energy change associated with the addition of an electron to a gaseous atom X(g) + e –  X – (g) electron affinity can be positive or negative

26. Why? ELECTRON AFFINITY general trend

28. 1. closed shells are energetically most stable 2. half-filled shells are energetically very stable DERIVATION OF ELECTRON AFFINITIES