Atomic Electron Configurations and Periodic Trends

1. Chapter 8 Atomic Electron Configurations and Chemical Periodicity

2. Chapter goals • Understanding the role magnetism plays in determining and revealing atomic structure. • Understand effective nuclear charge and its role in determining atomic properties. • Write the electron configuration of neutral atoms and monatomic ions. • Understand the fundamental physical properties of the elements and their periodic trends.

3. Electron Spin and the Fourth Quantum Number • The fourth quantum number is the spin quantum number which has the symbol ms. • The spin quantum number only has two possible values. ms = +1/2 or −1/2 ms = ± 1/2 • This quantum number tells us the spin and orientation of the magnetic field of the electrons. • Wolfgang Pauli discovered the Exclusion Principle in 1925. No two electrons in an atom can have the same set of 4 quantum numbers, n, l, ml, and ms

4. Electron Spin • Spin quantum number effects: • Every orbital can hold up to two electrons. • Consequence of the Pauli Exclusion Principle. • The two electrons are designated as having • one spin up  ms = +1/2 • and one spin down ms = −1/2 • Spin describes the direction of the electron’s magnetic field.

5. Paramagnetism and Diamagnetism • Unpaired electrons have their spins aligned  or  (in diff. orbitals) • This increases the magnetic field of the atom. Total spin  0, because they add up. • Atoms with unpaired electrons are called paramagnetic. • Paramagnetic atoms are attracted to a magnet.

6. Paramagnetism and Diamagnetism • Paired electrons have their spins unaligned . (in the same orbital) • Paired electrons have no net magnetic field. Total spin = 0, because of cancellation, ½ − ½ = 0 • Atoms with no unpaired electrons are called diamagnetic. • Diamagnetic atoms are not attracted to a magnet.

7. Atomic Orbitals, Spin, and # of Electrons • Because two electrons in the same orbital must be paired (due to Pauli’s Exclusion Principle), it is possible to calculate the number of orbitals and the number of electrons in each n shell. • The number of orbitals per n level is given by n2 (see table at end of chapter 7.) • The maximum number of electrons per n level is 2n2 (two electrons per orbital.) • The value is 2n2 because of the two paired electrons per orbital.

8. #orbitals Max n2 ml n shell l subshell #e– s 1 0 1 0 1 2 K s 2 0 2 0 1 L 8 4 p 6 –1,0,1 1 3 s 2 0 3 0 1 M 18 p 9 6 –1,0,1 1 3 d 10 2 5 -2,-1,0,1,2 s 0 4 0 1 N 2 6 p –1,0,1 1 3 32 16 10 d 2 5 -2,-1,0,1,2 -3,-2,-1,0,1,2,3 f 14 3 7

9. Atomic Subshell Energies and Electron Assignments • The principle that describes how the periodic chart is a function of electronic configurations is the Aufbau Principle. • The electron that distinguishes an element from the previous element enters the lowest energy atomic orbital available.

10. Penetrating and Shielding • the radial distribution function shows that the 2s orbital penetrates more deeply into the 1s orbital than does the 2p • the weaker penetration of the 2p sublevel means that electrons in the 2p sublevel experience more repulsive force, they are more shielded from the attractive force of the nucleus • the deeper penetration of the 2s electrons means electrons in the 2s sublevel experience a greater attractive force to the nucleus and are not shielded as effectively • the result is that the electrons in the 2s sublevel are lower in (more negative) energy than the electrons in the 2p

11. Atomic Subshell Energies and Electron Assignments The Aufbau Principle describes the electron filling order in atoms. This is product of the effective nuclear charge, Z*, Zeff For the same n, Z* is higher for s orbital: s > p > d > f Then, e− in s is the most attracted by nucleus and has the lowest energy

12. Atomic Subshell Energies and Electron Assignments One mnemonic to remember the correct filling order for electrons in atoms is the increasing (n + ) value

13. Atomic Subshell Energies and Electron Assignments or we can use this periodic chart

14. Atomic Electron Configurations • Now we will use the Aufbau Principle to determine the electronic configurations of the elements on the periodic chart. • 1st row elements

15. Atomic Electron Configurations Hund’srule tells us that the electrons will fill the p and d orbitals by placing electrons in each orbital singly and with same spin until half-filled. That is the rule of maximum spin. Then the electrons will pair to finish the p orbitals. Electrons in orbitals of or same kind, such as p or d orbitals, in the same shell (n), have the same energy; the are said to be degenerate.

16. Atomic Electron Configurations 3rd row elements…

17. Atomic Electron Configurations 4th row elements…

18. Atomic Electron Configurations 4th row elements…

19. Atomic Electron Configurations 4th row elements… The five d orbitals are degenerate

20. Atomic Electron Configurations 4th row elements…

21. Atomic Electron Configurations 4th row elements… The five d orbitals are degenerate

22. Atomic Electron Configurations 4th row elements…

23. Atomic Electron Configurations 4th row elements… The [Ar] 4s1 3d5 configuration of Cr is more stable than [Ar] 4s2 3d4 (expected)

24. Atomic Electron Configurations 4th row elements… The [Ar] 4s1 3d10full d configuration of Cu is more stable than [Ar] 4s2 3d9 (expected)

25. Atomic Electron Configurations 4th row elements…

26. Atomic Electron Configurations 4th row elements… (remember Hund’s rule):   __ is better (lower energy) than  __ __ 4p 4p

27. Atomic Electron Configurations Lanthanides (4f) 56Ba [Xe] 6s2 57La [Xe] 5d1 6s2 58Ce [Xe] 4f1 5d1 6s2 59Pr [Xe] 4f3 6s2 Praseodymium 70Yb [Xe] 4f14 6s2 Ytterbium 71Lu [Xe] 4f14 5d1 6s2 Lutetium

28. Periodic Table

29. s, p, d, and f-block in the Periodic Table

30. 1A 8A 2 1 2A 3A 7A 6A He H 5A 4A 6 8 5 9 10 7 3 4 N C O F Ne B Li Be 11 12 18 14 15 16 17 13 Al Si Ar Cl P Na S Mg 19 20 32 33 36 35 31 34 K Se Ca Br Ga Kr Ge As 37 38 49 50 53 51 54 52 In Sb Te I Sn Rb Sr Xe 55 56 82 81 83 85 86 84 Ba Rn Cs Pb Tl Bi At Po 87 88 Fr Ra P 1 2 (P–1)d 3 3B 4B 6B 8B 8B 8B 1B 5B 7B 2B 21 26 22 27 24 25 28 29 30 23 4 Sc Fe Mn Co Ni Zn Ti Cr Cu V 46 47 43 39 44 40 45 41 42 48 5 Ag Ru Tc Rh Cd Pd Y Zr Mo Nb 57 76 78 74 75 77 72 73 80 79 6 La Os Pt Ta Ir Au Hf W Hg Re 89 105 109 104 108 106 107 7 (P)p Ac Rf Db Sg Bh Hs Mt 63 69 68 71 60 70 59 61 64 62 66 (P)s 58 65 67 Lu Sm Gd Eu Ho Ce Tb Dy Er Pr Nd Pm Tm Yb 93 91 92 96 94 95 99 90 97 98 100 101 102 103 (P–2)f Th Lr U Np Cm Pa Pu Am Bk Cf Es Fm Md No

31. Valence Electrons electrons in shell with highest n, i.e., the outermost electrons, those beyond the core electrons 1s2 2s2 2p63s1 1s2 2s2 2p63s2 3p2 1s2 2s2 2p6 3s2 3p6 3d104s2 4p6 1s22s2 1s2 2s2 2p63s2 3p6 4s2 3d7 They determine the chemical properties of an element. For the representative elements, they are the ns and np electrons; for transition elements they are the ns and (n−1)d electrons.

32. 1A P 1 1s1 1 H 3 2s1 2 Li 11 3 3s1 Na # of valence electrons = 1 19 4 4s1 K 37 5 5s1 Rb 55 6s1 6 Cs 87 7s1 7 Fr

33. 2A 4 2s2 Be 12 3s2 Mg # of valence electrons = 2 20 4s2 Ca 38 5s2 Sr 56 6s2 Ba 88 7s2 Ra

34. 3A 5 2s2 2p1 B 13 3s2 3p1 Al # of valence electrons = 3 31 4s2 4p1 Ga 49 5s2 5p1 In 81 6s2 6p1 Tl

35. # of valence electrons = 7 7A 9 2s2 2p5 F 17 3s2 3p5 Cl 35 4s2 4p5 Br 53 5s2 5p5 I 85 6s2 6p5 At For the representative elements, the # of valence electrons = # of group

36. 1A 8A 2 1 2A 3A 7A 6A H He 5A 4A 6 8 5 9 10 7 3 4 N C O F Ne B Li Be 11 12 18 14 15 16 17 13 Al Si Ar Na Cl P S 3B 4B 6B 8B 8B 8B 1B Mg 5B 7B 2B 19 20 21 26 22 27 24 25 28 29 30 23 32 33 36 35 31 34 K Ca Sc Fe Mn Co Ni Se Br Ti Cr Cu Ga V Kr Ge As 46 37 47 38 43 39 44 40 45 41 42 49 50 53 48 51 54 52 Ag Ru Tc Rh Cd Pd Rb In Sb Te I Sr Y Zr Sn Mo Nb Xe 55 56 57 76 78 74 75 77 72 73 82 81 83 85 86 80 84 79 Ba Cs La Os Pt Rn Ta Ir Au Hf Pb Tl Bi At W Hg Re Po 87 88 89 105 109 104 108 106 107 Fr Ra The element X has the valence shell electron configuration, ns2 np4. X belongs to what group? chalcogens Zn Ac Unh Uns Uno Une Unp Unq

37. Energy (Orbital) Diagram 4p 3d 4s 3p 3s E 2p 2s Be 1s2 2s2 1s

38. Orbital Box Diagrams Be 1s 2s 3s 2p

39. Orbital Box Diagrams N 1s 2s 2p

40. Formation of Cations electrons lost from subshell with highest n and l first (from valence electrons) examples K 1s2 2s2 2p6 3s2 3p6 4s1 [Ar] 4s1 K+ 1s2 2s2 2p6 3s2 3p6 [Ar]

41. Ca 1s2 2s2 2p6 3s2 3p6 4s2 [Ar] 4s2 Ca2+ 1s2 2s2 2p6 3s2 3p6 [Ar] Al 1s2 2s2 2p63s2 3p1 Al3+ [Ne] In [Kr] 4d105s2 5p1 In3+ [Kr] 4d10

42. Transition Metal Cations In the process of ionization transition metals the ns electrons are lost before the (n-1)d Fe: [Ar] 3d6 4s2 Fe2+: [Ar] 3d6 Fe2+: [Ar] 3d6  Fe3+: [Ar] 3d5 Cu: [Ar] 3d10 4s1 Cu+: [Ar] 3d10 Cu+: [Ar] 3d10  Cu2+: [Ar] 3d9 Fe, Fe2+, Fe3+, Cu, and Cu2+ are paramagnetic

43. Two problems of ions, charge, and electron configuration An anion has a 3− charge and electron configuration 1s2 2s2 2p6 3s2 3p6. What is the symbol of the ion? The neutral atom has gained 3e- to form the ion, then the neutral atom had 15 e-. In the neutral atom the # e- = # p+ = Atomic number, that is 15. The element is, then, phosphorus (phosphorus). Symbol of ion is P3−. A cation has a 2+charge and its electron configuration is [Ar] 3d7. What is the symbol of the ion? Here, the neutral atom has lost 2e-. It is a transition metal, due to the 3d electrons. Remember they firstly lose e-s in 4s orbital. Symbol of ion is Co2+. Neutral atom has 18 + 7 + 2 = 27 e- = 27 p+ = atomic # [Ar] 3d7 lost

44. Atomic Properties and Periodic Trends Periodic Properties of the Elements • Atomic Radii • Ionization Energy • Electron Affinity • Ionic Radii

45. Atomic Properties and Periodic Trends • Establish a classification scheme of the elements based on their electron configurations. • Noble Gases • All of them have completely filled electron shells. They are not very reactive. • Since they have similar electronic structures, their chemical reactions are similar. • He 1s2 • Ne [He] 2s2 2p6 • Ar [Ne] 3s2 3p6 • Kr [Ar] 4s2 4p6 • Xe [Kr] 5s2 5p6 • Rn [Xe] 6s2 6p6

46. Atomic Properties and Periodic Trends Representative Elements are the elements in A groups on periodic chart. These elements will have their “last” electron in an outer s or p orbital. These elements have fairly regular variations in their properties. Metallic character, for expl, increases from right to left and top to bottom.

47. Atomic Properties and Periodic Trends • d-Transition Elements Elements on periodic chart in B groups. Sometimes called transition metals. • Each metal has d electrons. nsx (n-1)dy configurations • These elements make the transition from metals to nonmetals. • Exhibit smaller variations from row-to-row than the representative elements.

48. Atomic Properties and Periodic Trends • f - transition metals Sometimes called inner transition metals. • Electrons are being added to f orbitals. • Electrons are being added two shells below the valence shell! • Consequently, very slight variations of properties from one element to another.

49. Atomic Properties and Periodic Trends Outermost electrons (valence electrons) have the greatest Influence on the chemical properties of elements.

50. Atomic Properties and Periodic Trends Atomic radii describe the relative sizes of atoms. Atomic radii increase within a column going from the top to the bottom of the periodic table. The outermost electrons are assigned to orbitals with increasingly higher values of n. The underlying electrons require some space, so the electrons of the outer shells must be further from the nucleus.