1. Chapter Seven: Atomic Structure and Periodicity

2. Electrons in atoms are arranged as SHELLS (n) SUBSHELLS (l) ORBITALS (m l ) Arrangement of Electrons in Atoms

3. When n = 1, then l = 0 This shell has a single orbital (1s) to which 2 electrons can be assigned. When n = 2, then l = 0, 1 2s orbital 2 electrons 2s orbital 2 electrons three 2p orbitals6 electrons three 2p orbitals6 electrons TOTAL = 8 electrons Electrons in Atoms

4. The number of shells increases with the shell value (n). Therefore each higher shell holds more electrons. (more orbitals) Electrons in Shells

5. Electrons in an atom are arranged by: (quantum number) Principle energy levels(shells)n angular energy levels(sub-shells)l oriented energy levels(orbitals)mlml Within an individual orbital, there may only be two electrons differentiated by their spin. Spin states(electrons)msms Quantum Numbers

6. Based on theoretical and experimental studies of electron distributions in atoms, chemists have found there are two general rules that help predict these arrangements: 1. Electrons are assigned to subshells in order of increasing “n + l” value. 2. For two subshells with the same value of “n + l” electrons are assigned first to the subshell of lower n. Atomic Subshell Energies & Electron Assignments

7. In a Hydrogen atom (1–electron) the orbitals of a subshell are equal in energy (degenerate) 1s 2s3s4s2p3p3d Energy Single-Electron Atom Energy Levels

8. 1s 2s 3s 4s 2p 3p 3d E = n + l Screening results in the 4s- orbital having a lower energy that that of the 3d-orbital. 4s: n + l = 4 + 0 = 4 vs. 3d: n + l = 3 + 2 = 5 Multi-Electron Atom Energy Levels

9. Z* is the net charge experienced by a particular electron in a multi-electron atom resulting from a balance of the attractive force of the nucleus and the repulsive forces of other electrons. Z* increases across a period owing to incomplete screening by inner electrons This explains why E(4s electron) < E(3p electron) Z*  [ Z  (no. inner electrons) ] Charge felt by 2s electron in: LiZ* = 3  2 = 1 BeZ* = 4  2 = 2 B Z* = 5  2 = 3 and so on! Effective Nuclear Charge, Z*

10. Electron Configurations & the Periodic Table

11. Electron Filling Order

12. In the H-atom, all subshells of same n have same energy. In a multi-electron atom: 1.subshells increase in energy as value of n + l increases. 2.for subshells of same n + I, the subshell with lower n is lower in energy. Assigning Electrons to Subshells

13. Aufbau Principle: Lower energy orbitals fill first. Hund’s Rule: Degenerate orbitals (those of the same energy) are filled with electrons until all are half filled before pairing up of electrons can occur. Pauli exclusion principle: Individual orbitals only hold two electrons, and each should have different spin. “s” orbitals can hold 2 electrons “p” orbitals hold up to 6 electrons “d” orbitals can hold up to 10 electrons Orbital Filling Rules:

14. Hund’s Rule. Degenerate orbitals are filled with electrons until all are half-filled before pairing up of electrons can occur. Consider a set of 2p orbitals: 2p Electrons fill in this manner    Orbital Filling: The “Hund’s Rule”

15. “Pauli exclusion principle” Individual orbitals only hold two electrons, and each should have opposite spin. Consider a set of 2p orbitals: 2p Electrons fill in this manner       = spin up  = spin down *the convention is to write up, then down. Orbital Filling: The “Pauli Principle ”

16. Electrons fill the orbitals from lowest to highest energy. The electron configuration of an atom is the total sum of the electrons from lowest to highest shell. Example: Nitrogen:N has an atomic number of 7,therefore 7 electrons 1s2s2p      1s 2 2s 2 2p 3 Orbitals Electron Configuration (spdf) notation: Electron Configuration: Orbital Box Notation

17. Atomic Electron Configurations

18. Group 4A Atomic number = 6 6 total electrons 1s 2 2s 2 2p 2 1s2s2p      Carbon

19. Orbital box notation: 1s2s2p3s3p This corresponds to the energy level diagram: 1s 2s 2p 3s 3p Electron Configuration in the 3 rd Period

20. Orbital box notation: 1s2s2p3s3p Aluminum: Al (13 electrons) 1s 2s 2p 3s 3p                    1s 2 2s 2 2p 6 3s 2 spdf Electron Configuration 3p 1 Electron Configuration in the 3 rd Period

21. 3s3p [Ne] The electron configuration of an element can be represented as a function of the core electrons in terms of a noble gas and the valence electrons. Orbital Box Notation   Noble gas Notation [Ne] 3s 2 3p 2 Full electron configuration spdf notation 1s 2 2s 2 2p 6 3s 2 3p 2 Noble Gas Notation

22. The innermost electrons (core) can be represented by the full shell of noble gas electron configuration: 1s 2 2s 2 = [He], 1s 2 2s 2 2p 6 = [Ne], 1s 2 2s 2 2p 6 3s 2 3p 6 = [Ar]... The outermost electrons are referred to as the “Valence” electrons”. Element Full Electron Config. Noble Gas Notation Mg1s 2 2s 2 2p 6 3s 2 [Ne] 3s 2 Core or Noble Gas Notation

23. All 4th period and beyond d-block elements have the electron configuration [Ar] ns x (n - 1)d y Where n is the period and x, y are particular to the element. CopperIronChromium Transition Metal

24. Electron Configurations are written by shell even though the electrons fill by the periodic table: Ni: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 8 last electron to fill:3d 8 electron configuration by filling: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 8 electron configuration by shell: (write this way) Transition Elements

26. f-block elements: These elements have the configuration [core] ns x (n - 1)d y (n - 2)f z Where n is the period and x, y & z are particular to the element. Cerium: [Xe] 6s 2 5d 1 4f 1 Uranium: [Rn] 7s 2 6d 1 5f 3 Lanthanides & Actinides

27. © 2009, Prentice-Hall, Inc. Some Anomalies Some irregularities occur when there are enough electrons to half-fill s and d orbitals on a given row.

28. © 2009, Prentice-Hall, Inc. Some Anomalies For instance, the electron configuration for copper is [Ar] 4s 1 3d 10 rather than the expected [Ar] 4s 2 3d 9.

29. © 2009, Prentice-Hall, Inc. Some Anomalies This occurs because the 4s and 3d orbitals are very close in energy. These anomalies occur in f-block atoms, as well.

30. Filling Rules Aufbau Principle: – Electrons fill the lowest energy levels first. Pauli Exclusion Principle: – Electrons can fill two/orbital, providing they have opposite spins ( ) Hund’s Rule: – Orbitals of the same energy level (p,d,f) distribute electrons 1/orbital before adding the second.

31. Who are they? 1.1s 2 2s 2 2p 6 3s 2 2.[Kr]5s 2 4d 10 5p 5 3.1s 2 2s 2 2p 3 4.[Rn]7s 1 5.[Ar]4s 2 3d 10

32. Who are they? 1.1s 2 2s 2 2p 6 3s 2 Magnesium 2.[Kr]5s 2 4d 10 5p 5 Iodine 3.1s 2 2s 2 2p 3 Nitrogen 4.[Rn]7s 1 Francium 5.[Ar]4s 2 3d 10 Zinc

33. What’s wrong? Mg: [Ar]3s 2 Fe: 1s 2 2s 2 2p 6 3s 2 3d 6 4s 2 4d 6 Al: [Ne] 3s 3

34. Mistakes! Mg: [Ar]3s 2 Fe: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 4d 6 Al: [Ne] 3s 3

35. Corrections Mg: [Ne]3s 2 Fe: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 Al: [Ne] 3s 2 3p 1

36. Valence Electrons: Electrons at the highest energy level. Electrons available to be gained/lost/shared in a chemical reaction Valence electron configuration: ns x p x (total of 8 valence electrons) Representation of valence electrons in Lewis dot notation:

37. Formation of ions & Valence electrons Ions are formed when atoms either: 1. Cation (+): Give up (lose) electrons 2.Anion (-): Gain electrons Overall goal: stable noble gas notation Elements with < 4 valence electrons- form cations Elements with > 4 valence electrons for anions. Non-metals with 4 valence electrons- do not form ions Noble gases (8 valence electrons) – are unreactive

38. Ions & Electron Configuration Atoms or groups of atoms that carry a charge Cations- positive charge – Formed when an atom loses electron(s) – Na  Na+ + e- – Na+ : 1s22s22p6 (10 e-) = [Ne] Anions –negative charge – Formed when an atom gain electrons(s) – F + e-  F- – Cl- : 1s22s22p6 (10 e-) = [Ne]

39. Isoelectronic series: Ions that contain the same number of electrons. Example: Al3+, Mg 2+, Na +, F -, O 2-, N 3-

40. Transition Metals & Ions Transition metals: multi-valent ions (more than one possible charge) Electrons are removed from the highest quantum number first: Example: Cu + & Cu 2+ Cu + : [Ar] 3d 10 Cu 2+ : [Ar] 3d 9

41. Diamagnetic SubstancesDiamagnetic Substances: Are NOT attracted to a magnetic field Paramagnetic SubstancesParamagnetic Substances: ARE attracted to a magnetic field. unpaired electronsSubstances with unpaired electrons are paramagnetic. Electron Spin & Magnetism

42. Ions with UNPAIRED ELECTRONS are PARAMAGNETIC (attracted to a magnetic field). Ions without UNPAIRED ELECTRONS are DIAMAGNETIC (not attracted to a magnetic field). Fe 3+ ions in Fe 2 O 3 have 5 unpaired electrons. This makes the sample paramagnetic. Electron Configurations of Ions

43. Practice: Electron Configuration Write the following electron configurations 1. Silicon- orbital notation 2.Strontium – long (spdf) notation 3.Bismuth – noble gas notation 4.Silver – noble gas notation 5.O 2- - orbital notation 6.Fe 2+ - noble gas notation

44. Practice #2: Electron Configuration Write the following electron configurations 1. Chromium- long (spdf) notation 2.Sulfur – orbital notation 3.Ag+ – noble gas notation 4.Americium – noble gas notation 5.Mg 2+ - orbital notation 6.Krypton – noble gas notation

45. © 2009, Prentice-Hall, Inc. Development of Periodic Table Elements in the same group generally have similar chemical properties. Physical properties are not necessarily similar, however.

46. © 2009, Prentice-Hall, Inc. Periodic Trends In this chapter, we will rationalize observed trends in – Sizes of atoms and ions. – Ionization energy. – Electron affinity.

47. Atomic and ionic size Ionization energy Electron affinity Higher effective nuclear charge Electrons held more tightly Larger orbitals. Electrons held less tightly. General Periodic Trends

48. © 2009, Prentice-Hall, Inc. Effective Nuclear Charge In a many-electron atom, electrons are both attracted to the nucleus and repelled by other electrons. The nuclear charge that an electron experiences depends on both factors.

49. © 2009, Prentice-Hall, Inc. Effective Nuclear Charge The effective nuclear charge, Z eff, is found this way: Z eff = Z − S where Z is the atomic number and S is a screening constant, usually close to the number of inner electrons.

50. © 2009, Prentice-Hall, Inc. What Is the Size of an Atom? The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei.

51. © 2009, Prentice-Hall, Inc. Sizes of Atoms Bonding atomic radius tends to… …decrease from left to right across a row (due to increasing Z eff ). …increase from top to bottom of a column (due to increasing value of n).

52. © 2009, Prentice-Hall, Inc. Ionization Energy The ionization energy is the amount of energy required to remove an electron from the ground state of a gaseous atom or ion. – The first ionization energy is that energy required to remove first electron. – The second ionization energy is that energy required to remove second electron, etc.

53. © 2009, Prentice-Hall, Inc. Ionization Energy It requires more energy to remove each successive electron. When all valence electrons have been removed, the ionization energy takes a quantum leap.

54. © 2009, Prentice-Hall, Inc. Trends in First Ionization Energies As one goes down a column, less energy is required to remove the first electron. – For atoms in the same group, Z eff is essentially the same, but the valence electrons are farther from the nucleus.

55. © 2009, Prentice-Hall, Inc. Trends in First Ionization Energies Generally, as one goes across a row, it gets harder to remove an electron. – As you go from left to right, Z eff increases.

56. © 2009, Prentice-Hall, Inc. Trends in First Ionization Energies However, there are two apparent discontinuities in this trend.

57. © 2009, Prentice-Hall, Inc. Trends in First Ionization Energies The first occurs between Groups IIA and IIIA. In this case the electron is removed from a p- orbital rather than an s- orbital. – The electron removed is farther from nucleus. – There is also a small amount of repulsion by the s electrons.

58. © 2009, Prentice-Hall, Inc. Electron Affinity Electron affinity is the energy change accompanying the addition of an electron to a gaseous atom: Cl + e −  Cl −

59. © 2009, Prentice-Hall, Inc. Trends in Electron Affinity In general, electron affinity becomes more exothermic as you go from left to right across a row.

60. © 2009, Prentice-Hall, Inc. Trends in Electron Affinity There are again, however, two discontinuities in this trend.

61. © 2009, Prentice-Hall, Inc. Trends in Electron Affinity The first occurs between Groups IA and IIA. – The added electron must go in a p-orbital, not an s-orbital. – The electron is farther from nucleus and feels repulsion from the s- electrons.

62. © 2009, Prentice-Hall, Inc. Trends in Electron Affinity The second occurs between Groups IVA and VA. – Group VA has no empty orbitals. – The extra electron must go into an already occupied orbital, creating repulsion.

63. © 2009, Prentice-Hall, Inc. Sizes of Ions Ionic size depends upon: – The nuclear charge. – The number of electrons. – The orbitals in which electrons reside.

64. © 2009, Prentice-Hall, Inc. Sizes of Ions Cations are smaller than their parent atoms. – The outermost electron is removed and repulsions between electrons are reduced.

65. © 2009, Prentice-Hall, Inc. Sizes of Ions Anions are larger than their parent atoms. – Electrons are added and repulsions between electrons are increased.

66. © 2009, Prentice-Hall, Inc. Sizes of Ions Ions increase in size as you go down a column. – This is due to increasing value of n.

67. © 2009, Prentice-Hall, Inc. Sizes of Ions In an isoelectronic series, ions have the same number of electrons. Ionic size decreases with an increasing nuclear charge.

68. Problem: Rank the following ions in order of decreasing size? Na +,N 3-,Mg 2+,F–F– O 2–

69. Problem: Rank the following ions in order of decreasing size? Na +,N 3-,Mg 2+,F–F– O 2– Ion# of protons# of electronsratio of e/p Na + N 3- Mg 2+ F–F– O 2–

70. Problem: Rank the following ions in order of decreasing size? Na +,N 3-,Mg 2+,F–F– O 2– Ion# of protons# of electronsratio of e/p Na + N 3- Mg 2+ F–F– O 2– 11 7 12 9 8

71. Problem: Rank the following ions in order of decreasing size? Na +,N 3-,Mg 2+,F–F– O 2– Ion# of protons# of electronsratio of e/p Na + N 3- Mg 2+ F–F– O 2– 11 7 12 9 8 10

72. Problem: Rank the following ions in order of decreasing size? Na +,N 3-,Mg 2+,F–F– O 2– Ion# of protons# of electronsratio of e/p Na + N 3- Mg 2+ F–F– O 2– 11 7 12 9 8 10 0.909 1.43 0.833 1.11 1.25

73. Na + N 3- Mg 2+ F–F– O 2– Ion 0.909 1.43 0.833 1.11 1.25 e/p ratio Since N 3- has the highest ratio of electrons to protons, it must have the largest radius.

74. Na + N 3- Mg 2+ F–F– O 2– Ion 0.909 1.43 0.833 1.11 1.25 e/p ratio Since N 3- has the highest ratio of electrons to protons, it must have the largest radius. Since Mg 2+ has the lowest, it must have the smallest radius. The rest can be ranked by ratio.

75. Na + N 3- Mg 2+ F–F– O 2– Ion 0.909 1.43 0.833 1.11 1.25 e/p ratio Since N 3- has the highest ratio of electrons to protons, it must have the largest radius. Since Mg 2+ has the lowest, it must have the smallest radius. The rest can be ranked by ratio. N 3- >O 2– >F–F– >Na + >Mg 2+ Decreasing size isoelectronic Notice that they all have 10 electrons: They are isoelectronic (same electron configuration) as Ne.

76. Moving through the periodic table: Atomic radii Ionization Energy Electron Affinity Down a groupIncreaseDecrease Becomes less exothermic Across a PeriodDecreaseIncrease Becomes more exothermic Summary of Periodic Trends