Atomic Electron Configurations and Chemical Periodicity

1. Atomic Electron Configurations and Chemical Periodicity Goals: Understand 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 for elements and monoatomic ions. Understand the fundamental physical properties of the elements and their periodic trends.

2. Arrangement of Electronsin Atoms • Electrons in atoms are arranged as: • Shells (n) • Subshells (l) • Orbitals (ml) • Electrons have _____. • ms, __________________ 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. • Substances 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, ml =0 • ms = +1/2 • He (2e) n=1, l=0, ml = 0, ms = +1/2 • n=1, l=0, ml = 0, ms = -1/2 1s 1s

8. Electrons in Atoms

9. A 2p electron can be designated by which set of quantum numbers? n l ml ms a. 1 0 0 +1/2 b. 2 1 0 +1/2 c. 2 2 +1 -1/2 d. 3 1 +2 +1/2 e. 3 2 +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. • For H atoms, E = - C(1/n2). E depends only on _____. • For many-electron atoms, energy depends on both ____ and _____.

11. Assigning Electrons to Subshells • 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. 1 e- atom

12. Electron cloud for 1s electrons 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

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. for H, atomic number = 1 Orbital Box notation spdf notation 1 no. of s 1 electrons 1s label of l value of n

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

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

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

19. 3p 3s 2p 2s 1s Carbon Group 4A Atomic number = 6 ___________ ---> 6 total electrons 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.

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

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

23. Sodium and Potassium –Noble gas notation Na - Group 1A Atomic number = 11 1s2 2s2 2p6 3s1 or “neon core” + 3s1 [Ne] 3s1 (uses rare gas notation) Note that we have begun a new period (3rd ) All Group 1A elements have [core]ns1 configurations, n=period number K – Atomic number = 19, 4th period ___________

24. 3p 3s 2p 2s 1s Phosphorus Group 5A Atomic number = 15 spdf: ______________ short: __________ All Group 5A elements have [core ] ns2 np3 configurations where n is the period number.

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] nsx (n - 1)dy and so are d-block elements. Chromium Iron 26e- Copper

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] nsx (n - 1)dy (n - 2)fz and so are f-block elements. Cerium [Xe] 6s2 5d1 4f1 Uranium [Rn] 7s2 6d1 5f3

30. Ion Configurations To form cations from elements remove 1 or more e- from subshell of highest n [or highest (n + l)]. P ---> P3+ [Ne] 3s2 3p3 - 3e- [Ne] 3s2 3p0 3p 3p 3s 3s 2p 2p 2s 2s 1s 1s

31. 2+ Fe Fe 4s 4s 3d 3d 3+ Fe 4s 3d Ion Configurations For transition metals, remove ns electrons and then (n - 1) electrons. Fe [Ar] 4s2 3d6 loses 2 electrons ---> Fe2+ [Ar] 4s0 3d6 To form cations, always remove electrons of highest n value first!

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 CO2 binds to hemoglobin and is released in the lungs. About 70% is converted by Carbonic Anhydrase into HCO3- ion, which remains in the blood plasma until the reverse reaction releases CO2 into the lungs. • Carbonic Anhydrase catalyzes the reversible hydration of CO2 to form bicarbonate anion and a proton:CO2 + H2O <==> HCO3- + H+ • Toxic metals like Cd2+ replace Zn2+ inactivating the enzyme.

34. PERIODIC TRENDS

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

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

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

41. Ion Sizes

42. Redox Reactions Why do metals lose electrons in their reactions? Why does Mg form Mg2+ ions and not Mg3+? 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 ---> Mg2+ (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 ---> Mg2+ (g) + e- Mg2+ (g) + ---> Mg3+ (g) + e- 7733 kJ Energy cost is very high to dip into a shell of lower n. This is why oxidation number = Group number.

46. 1- 11+ Effective nuclear charge • Sodium 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. valance 10- Effect = 11-10 = +1 core

47. Effective nuclear charge [Ne] core Radial electron density 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. Increase in Z* Atomic size Size decreases across a period owing to increase in Z*. Each added electron feels a greater and greater + charge. Large Small

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

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