1. Chapter 8 Periodic Properties of the Elements

2. Energy of atomic orbitals For an atom, electrons are in atomic orbitals.

3. Orbital Energy Levels for the Hydrogen Atom H atom: E only depends on ndegenerate

4. E depends on n and l same n, l ↑ ↔ E↑

6. A Picture of the Spinning Electron

8. Spin quantum number m s m s = +1/2 or −1/2 4 quantum numbers are used to specify an electron. How do electrons fill up atomic orbitals?

9. Pauli Exclusion Principle In a given atom, no two electrons can have the same set of four quantum numbers. An orbital can hold only two electrons, they must have opposite spins.

11. 1s 1 2s 1 2p 1 electron configuration Lowest energy: ground state Excited states ↑ 1s ↑ 2s ↑ 2p H atom orbital diagram

12. Now we can write the ground state electron configurations and draw orbital diagrams according to Pauli principle.

15. For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized. Hund’s rule Valence electrons: electrons in the outermost shell. involved in bonding Core electrons: inner electrons

17. Elements in the same group have similar valence electron configuration — similar chemical properties. Noble gases have 8 (He 2) valence electrons. Stable structure. Number of valence electrons = main group number Metals: tend to lose valence electrons to reach 8(2) valence electron. Nonmetals: tend to gain electrons to reach 8(2) valence electrons. Number of filled shells = period number

20. Periodic trends in atomic properties Atomic radius

21. Atomic Radii (in Picometers) for Selected Atoms

22. Atomic radius In a period: decreases from left to right In a group: increases from top to bottom

23. On the basis of periodic trends, choose the larger atom in each pair (if possible). Explain your choices. (a) N or F (b) C or Ge (c) N or Al (d) Al or Ge EXAMPLE 8.5 Atomic Size

24. Choose the larger atom or ion from each pair. (a) S or S 2– (b) Ca or Ca 2+ (c) Br – or Kr EXAMPLE 8.7 Ion Size

25. Periodic trends in atomic properties Atomic radius Ionization energy

26. Energy required to remove an electron from a gaseous atom or ion. X(g)  X + (g) + e − X + (g)  X 2+ (g) + e − Ionization energy first ionization energy second ionization energy

28. Ionization energy In a period: increases from left to right In a group: decreases from top to bottom (general trend)

29. Periodic trends in atomic properties Atomic radius Ionization energy Electron affinity

30. Energy change associated with the addition of an electron to a gaseous atom. X(g) + e −  X − (g) Electron affinity X(g) + e − X − (g) E EiEi EfEf ∆E = E f − E i = EA < 0

32. Electron affinity In a period: increases from left to right In a group: no clear trend (very rough trend)

33. Periodic trends in atomic properties Atomic radius Ionization energy Electron affinity Remember the trends

34. 1. Arrange the following groups of atoms in order of increasing size. a)Te, S, Se; b) K, Br, Ni; c) Ba, Si, F 2. Arrange the atoms in Ex. 1 in order of increasing first ionization energy.

35. 4, 7, 43, 45, 51, 53, 55, 61, 63, 65, 67, 69, 73, 75 Chapter 8 Problems