1. Atomic Electron Configurations and Chemical Periodicity Goals: 1.Understand the role magnetism plays in determining and revealing atomic structure. 2.Understand effective nuclear charge and its role in determining atomic properties. 3.Write the electron configuration for elements and monoatomic ions. 4.Understand the fundamental physical properties of the elements and their periodic trends.

2. Arrangement of Electrons in Atoms Electrons in atoms are arranged as: Shells (n) Subshells ( l ) Orbitals (m l ) Electrons have _____. m s, __________________ quantum number, = +1/2 and -1/2 Complete description of electrons requires _______ quantum numbers.

3. Electron Spin Magnetic Quantum Number

4. Electron Spin and Magnetism ____________: NOT attracted to a magnetic field ___________: substance is attracted to a magnetic field. unpaired electronsSubstances with unpaired electrons are ______________.

5. Electron Spin and Magnetism H atoms, each has a single electron, they are paramagnetic – when an external magnetic field is applied, the electron magnets align with the field. He atoms, with two electrons, are diamagnetic. –We assumed opposite spin orientations – spins are __________.

6. The Pauli Exclusion Principle No two electrons in an atom can have the same set of four quantum numbers. Therefore, Each orbital can be assigned no more than ____ electrons!

7. Orbital Box Diagrams When n = 1, then l = 0 this shell has a single orbital (1s) to which 2e- can be assigned. H(1e)n=1, l =0, m l =0 m s = +1/2 He (2e) n=1, l =0, m l = 0, m s = +1/2 n=1, l =0, m l = 0, m s = -1/2 1s

8. Electrons in Atoms

9. A 2p electron can be designated by which set of quantum numbers? n l m l m s a.100+1/2 b.210+1/2 c.22+1-1/2 d.31+2+1/2 e.32+1+1/2 Students should be familiar with the values and meaning of quantum numbers.

10. Electrons in Atoms Electrons generally assigned to orbitals of successively higher energy.Electrons generally assigned to orbitals of successively higher energy. For H atoms, E = - C(1/n 2 ). E depends only on _____.For H atoms, E = - C(1/n 2 ). E depends only on _____. For many-electron atoms, energy depends on both ____ and _____.For many-electron atoms, energy depends on both ____ and _____.

11. Assigning Electrons to Subshells 1 e- atom In many-electron atom:In many-electron atom: a) subshells increase in energy as value of _______increases. b) for subshells of same _____, subshell with lower n is lower in energy.

12. Effective Nuclear Charge, Z* Z* - the nuclear charge experienced by a particular electron in a multielectron atom, as modified by the presence of the other electrons. Li has 3 p (+) and 3 e (-) 2 e in 1 s orbital ; 1 e in 2 s orbital e- in 2s should “see” a +1 charge, but it sees 1.28 C has 6 p (+) and 6 e (-) 2 e- in 1s ; 2 e- in 2s ; 2 e- in 2p e- in 2s should “see” +3, but see 3.22 e- in 2p should “see” a +2 charge, but see 3.14 Electron cloud for 1s electrons

13. Effective Nuclear Charge, Z* Z* is _______ for s electrons than for p electrons. – s electrons always have a lower energy than p electrons in the same quantum shell. The Z* _________ across a period. The 2s electron PENETRATES the region occupied by the 1s electron. –2s electron experiences a _________ positive charge than expected.

14. Atomic Electron Configurations The arrangements of __________ in the elements in the ground state. In general, electrons are assigned to orbitals in order of increasing ________. Electron configuration can be given with the orbital box diagram, or with the spdf notation. 1 1 s value of n label of l no. of electrons spdf notation for H, atomic number = 1 Orbital Box notation 1s

16. Electron Configurations The outermost electrons of an element are assigned to the indicated orbitals.

17. Lithium Group 1A Atomic number = 3 ________ ---> 3 total electrons 1s 2s 3s 3p 2p

18. Boron Group 3A Atomic number = 5 ___________ ---> 5 total electrons 5 total electrons 1s 2s 3s 3p 2p

19. Carbon Here we see for the first time ___________ RULE: When placing electrons in a set of orbitals having the same energy, we place them singly as long as possible. Group 4A Atomic number = 6 ___________ ---> 6 total electrons 6 total electrons 1s 2s 3s 3p 2p

20. Hund’s Rule The most stable arrangement of electrons is that with the __________ _____________________, all with the same ________ direction. This arrangement makes the total energy of an atom as low as possible.

21. Nitrogen and Oxygen Group 5A Atomic number = 7 _________---> 7 total electrons 7 total electrons 1s 2s 3s 3p 2p Group 6A 1s 2s 3s 3p 2p

22. 1s 2s 3s 3p 2p Fluorine and Neon Note that we have reached the end of the 2nd period, and the 2nd shell is ________!Note that we have reached the end of the 2nd period, and the 2nd shell is ________! Group 7A Atomic number = 9 ____________---> 9 total electrons 9 total electrons Group 8A Atomic number = 10 _____________ ---> 10 total electrons 10 total electrons 1s 2s 3s 3p 2p

23. Sodium and Potassium – Noble gas notation Na - Group 1A Atomic number = 11 1s 2 2s 2 2p 6 3s 1 or “neon core” + 3s 1 [Ne] 3s 1 (uses rare gas notation) Note that we have begun a new period (3 rd ) All Group 1A elements have [core]ns 1 configurations, n=period number K – Atomic number = 19, 4 th period ___________

24. Phosphorus All Group 5A elements have [core ] ns 2 np 3 configurations where n is the period number. Group 5A Atomic number = 15 spdf: ______________ short: __________ 1s 2s 3s 3p 2p

25. Transition Metals 3d orbitals used for Sc- Zn (Table 8.4) and so are d-block elements.

26. Transition Metals All 4th period elements have the configuration [argon] ns x (n - 1)d y and so are d-block elements. Copper Iron Chromium 26e-

27. Lantanides and Actinides 4f orbitals used for Ce - Lu and 5f for Th - Lr (Table 8.2) and so are f-block elements.

28. Lantanides and Actinides All these elements have the configuration [core] ns x (n - 1)d y (n - 2)f z and so are f-block elements. Cerium [Xe] 6s 2 5d 1 4f 1 Uranium [Rn] 7s 2 6d 1 5f 3

30. Ion Configurations To form cations from elements remove 1 or more e- from subshell of highest n [or highest (n + l)]. P ---> P 3+ [Ne] 3s 2 3p 3 - 3e- [Ne] 3s 2 3p 0 1s 2s 3s 3p 2p 1s 2s 3s 3p 2p

31. Ion Configurations For transition metals, remove ns electrons and then (n - 1) electrons. Fe [Ar] 4s 2 3d 6 loses 2 electrons ---> Fe 2+ [Ar] 4s 0 3d 6 To form cations, always remove electrons of highest n value first ! 3d 4s Fe 3+ 4s 3d 4s Fe 2+

32. Practice Which substance will be paramagnetic? V +5 or Fe +3 Students should be familiar with writing electronic configurations and identifying diamagnetic vs. paramagnetic materials.

33. Ion Size Makes a BIG Difference About 20% of the CO 2 binds to hemoglobin and is released in the lungs. About 70% is converted by Carbonic Anhydrase into HCO 3- ion, which remains in the blood plasma until the reverse reaction releases CO 2 into the lungs. Carbonic Anhydrase catalyzes the reversible hydration of CO 2 to form bicarbonate anion and a proton: CO 2 + H 2 O HCO 3- + H+ Toxic metals like Cd 2+ replace Zn 2+ inactivating the enzyme.

34. PERIODIC TRENDS

35. Periodic Trends Atomic and ionic sizeAtomic and ionic size Ionization energyIonization energy Electron affinityElectron affinity Higher effective nuclear charge Electrons held ______ tightly Larger orbitals. Electrons held ____ tightly.

36. Atomic size Size goes ____ on going down a group. See Figure 8.9. Because electrons are added further from the nucleus, there is ______ attraction. Size goes _______ on going across a period.

37. Effective Nuclear Charge, Z* AtomZ* Experienced by Electrons in Valence Orbitals (Outermost) Li+1.28 Be------- B+2.58 C+3.22 N+3.85 O+4.49 F+5.13 Increase in Z* across a period [Values calculated using Slater’s Rules]

38. Atomic size- Transition Metals 3d subshell is inside the 4s subshell. 4s electrons feel a more or less constant Z*. Sizes stay about the same and chemistries are similar!

39. Ion Sizes CATIONS are __________ than the atoms from which they come.CATIONS are __________ than the atoms from which they come. The electron/proton attraction has gone ______ and so size ____________.The electron/proton attraction has gone ______ and so size ____________. Li,152 pm 3e and 3p Li +, 78 pm 2e and 3 p + Forming a cation.

40. Ion Sizes ANIONS are __________ than the atoms from which they come.ANIONS are __________ than the atoms from which they come. The electron/proton attraction has gone ______and so size __________.The electron/proton attraction has gone ______and so size __________. Trends in ion sizes are the same as atom sizes.Trends in ion sizes are the same as atom sizes. Forming an anion. F, 71 pm 9e and 9p F -, 133 pm 10 e and 9 p -

41. Ion Sizes

42. Redox Reactions Why do metals lose electrons in their reactions? Why does Mg form Mg 2+ ions and not Mg 3+ ? Why do nonmetals take on electrons? Why do metals lose electrons in their reactions? Why does Mg form Mg 2+ ions and not Mg 3+ ? Why do nonmetals take on electrons?

43. Ionization Energy IE = energy required to remove an electron from an atom in the gas phase. Mg (g) + 738 kJ ---> Mg + (g) + e-

44. Ionization Energy Mg (g) + 738 kJ ---> Mg + (g) + e- Mg + (g) + 1451 kJ ---> Mg 2+ (g) + e- Mg + has 12 protons and only 11 electrons. Therefore, IE for Mg + > Mg.

45. Ionization Energy Mg (g) + 735 kJ ---> Mg + (g) + e- Mg + (g) + 1451 kJ ---> Mg 2+ (g) + e- Mg 2+ (g) + ---> Mg 3+ (g) + e- Energy cost is very high to dip into a shell of lower n. This is why oxidation number = Group number. 7733 kJ

46. Effective nuclear charge Sodium 11+ 10- 1- A valance electron in an atom is attracted to the nucleus of the atom and it is repelled by the other electrons in the atom: inner e- shield or screen the outer electrons from attraction of the nucleus. Effect = 11-10 = +1 core valance

47. Effective nuclear charge Radial electron density [Ne] core 3s For the 3s e- (valance e- of Na) there is a probability of being found close to the nucleus – there is a probability of experiencing a greater attraction than suggested. Zeff = +2.5 Electrons in 3s orbitals has a higher Zeff than 3p orbitals: subshells energy trend is: ns < np < nd

48. Atomic size Size decreases across a period owing to increase in Z*. Each added electron feels a greater and greater + charge. Large Small Increase in Z*

49. Ionization Energy IE ___________ across a period and ___________ down a group.

50. Ionization Energy As Z* increases, orbital energies “drop” and IE increases.

51. Trends in Ionization Energy IE increases across a period because Z* increases.IE increases across a period because Z* increases. –Metals lose electrons more easily than nonmetals. –Metals are good reducing agents. –Nonmetals lose electrons with difficulty. IE decreases down a group.IE decreases down a group. –Because size increases. –Reducing ability generally increases down the periodic table (easier to give e-). –See reactions of Li, Na, K

52. Ionization Energy

53. Electron Affinity A few elements GAIN electrons to form anions. Electron affinity is the energy change that occurs when _______________ to ___________________________. A(g) + e- ---> A - (g) Electron Affinity = ∆E

54. Electron Affinity

55. Affinity for electron _________ across a period (EA becomes more negative). Affinity __________ down a group (EA becomes less negative).

56. Summary

57. Practice Which of the following elements has the greater difference between its first and second ionization energies: C, Li, N, Be? Which should be smaller: the sulfide ion, S 2-, or a sulfur atom, S? Students should be familiar with periodic trends – IE, EA, Atomic size.

58. Remember Go over all the contents of your textbook. Practice with examples and with problems at the end of the chapter. Practice with OWL tutor. Work on your assignment for Chapter 8. Practice with the quiz on CD of Chemistry Now.