chapter7

chapter7 - 1 ATOMIC ELECTRON ATOMIC CONFIGURATIONS AND...

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Unformatted text preview: 1 ATOMIC ELECTRON ATOMIC CONFIGURATIONS AND PERIODICITY CONFIGURATIONS 2 Arrangement of Electrons Arrangement in Atoms in Electrons in atoms are arranged as SHELLS (n) SUBSHELLS (l) ORBITALS (ml) 3 Arrangement of Electrons Arrangement in Atoms in Each orbital can be assigned no Each more than 2 electrons! more This is tied to the existence of a 4th This quantum number, the electron spin quantum number, ms. spin 4 Electron Electron Spin Quantum Number, ms Can be proved experimentally that electron Can has an intrinsic property referred to as “spin.” Two spin directions are given by Two ms where ms = +1/2 and -1/2. 5 Now there are four! n shell shell 1, 2, 3, 4, ... l subshell 0, 1, 2, ... n - 1 ml orbital - l ... 0 ... + l ms electron spin +1/2 and -1/2 electron 6 Pauli Exclusion Principle No two electrons in the No same atom can have the same set of 4 quantum numbers. quantum That is, each electron has a That unique address. unique 7 Electrons in Atoms When n = 1, then l = 0 When this shell has a single orbital (1s) to which 2ecan be assigned. can When n = 2, then l = 0, 1 When 2s orbital 2e2e- three 2p orbitals three 6e- TOTAL = 8e8e- 8 Electrons in Atoms When n = 3, then l = 0, 1, 2 When 3s orbital three 3p orbitals three five 3d orbitals five TOTAL = 2e2e6e10e18e18e- 9 Electrons in Atoms When n = 4, then l = 0, 1, 2, 3 When 4s orbital 2e2ethree 4p orbitals 6ethree five 4d orbitals 10efive seven 4f orbitals 14eseven TOTAL = 32e32e- And many more! 10 11 Assigning Electrons to Atoms • Electrons generally assigned to orbitals of Electrons successively higher energy. successively • For H atoms, E = - C(1/n2). E depends only For atoms ). on n. on • For many-electron atoms, energy depends For many-electron energy on both n and l. • See Active Figure 7.1 and Figure 7.2 12 Assigning Electrons to Subshells • In H atom all subshells In of same n have same energy. energy. • In many-electron atom: a) subshells increase in a) energy as value of n + l increases. increases. b) for subshells of same n b) + l, subshell with lower subshell n is lower in energy. is 13 Electron Electron Filling Order Order See Figure 7.2 14 Effective Nuclear Charge, Z* • Z* is the nuclear charge experienced by Z* the outermost electrons. See Figure 7.3 See the • Explains why E(2s) < E(2p) • Z* increases across a period owing to Z* incomplete shielding by inner electrons. incomplete • Estimate Z* = [ Z - (no. inner electrons) ] Estimate • Charge felt by 2s e- in Li Z* = 3 - 2 = 1 Charge • Be Z* = 4 - 2 = 2 Be Z* •B Z* = 5 - 2 = 3 and so on! Z* 15 Effective Effective Nuclear Charge Charge See Figure 7.3 Z* is the nuclear Z* charge experienced by the outermost electrons. electrons. Electron cloud for 1s electrons 16 Writing Atomic Electron Writing Configurations Configurations Two ways of Two writing configs. One is called the spdf spdf notation for H, atomic number = 1 notation. notation. 1 value of n no. of electrons 17 Writing Atomic Electron Writing Configurations Configurations Two ways of Two writing configs. Other is called the orbital box notation. notation. ORBITAL BOX NOTATION for He, atomic number = 2 Arrows 2 depict electron spin 1s 1s One electron has n = 1, l = 0, ml = 0, ms = + 1/2 One Other electron has n = 1, l = 0, ml = 0, ms = - 1/2 Other 18 19 Electron Configurations Electron and the Periodic Table and See Active Figure 7.4 20 Lithium Group 1A Atomic number = 3 1s22s1 3 total electrons 3p 3s 2p 2s 1s 21 Beryllium 3p 3s 2p 2s 1s Group 2A Atomic number = 4 1s22s2 4 total electrons electrons 22 Boron 3p 3s 2p 2s 1s Group 3A Atomic number = 5 1s2 2s2 2p1 2s 2p 5 total electrons total 23 Carbon Group 4A Atomic number = 6 1s2 2s2 2p2 2s 2p 6 total electrons total 3p 3s 2p 2s 1s Here we see for the first time Here HUND’SRULE. When placing HUND When electrons in a set of orbitals having the same energy, we place them singly as long as possible. possible. 24 Nitrogen 3p 3s 2p 2s 1s Group 5A Atomic number = 7 1s2 2s2 2p3 2s 2p 7 total electrons total 25 Oxygen 3p 3s 2p 2s 1s Group 6A Atomic number = 8 1s2 2s2 2p4 2s 2p 8 total electrons total 26 Fluorine 3p 3s 2p 2s 1s Group 7A Atomic number = 9 1s2 2s2 2p5 2s 2p 9 total electrons total 27 Neon Group 8A Atomic number = 10 1s2 2s2 2p6 2s 2p 10 total electrons 10 3p 3s 2p 2s 1s Notethat wehavere d ache Note t hee of the2nd pe nd riod, and the2nd she is full! ll and 28 Electron Configurations of pBlock Elements 29 Sodium Group 1A Atomic number = 11 1s2 2s2 2p6 3s1 or 2s 2p 3s “neon core” + 3s1 [Ne] 3s1 (uses rare gas notation) (uses Note that we have begun a new period. All Group 1A elements have All [core]ns1 configurations. [core]ns 30 Aluminum Group 3A Atomic number = 13 1s2 2s2 2p6 3s2 3p1 2s 2p 3s 3p [Ne] 3s2 3p1 [Ne] 3p 3p All Group 3A e m nts have[core le e ns2 np1 configurations whe n is re thepe num r. be t he riod 3s 2p 2s 1s 31 Phosphorus Yellow P Group 5A Atomic number = 15 1s2 2s2 2p6 3s2 3p3 2s 2p 3s 3p [Ne] 3s2 3p3 [Ne] 3p Red P 3p All Group 5A e m nts have le e All [core] ns2 np3 configurations whe n is thepe num r. re riod be whe 3s 2p 2s 1s 32 Calcium Group 2A Atomic number = 20 1s2 2s2 2p6 3s2 3p6 4s2 2s 2p 3s 3p [Ar] 4s2 [Ar] All Group 2A elements have [core]ns2 configurations where n [core]ns is the period number. is 33 Electron Configurations Electron and the Periodic Table and 34 Transition Metals Table 7.4 All 4th period elements have the All configuration [argon] nsx (n - 1)dy [argon] and so are d-block elements. d-block Chromium Iron Copper Transition Element Transition Configurations Configurations 3d orbitals used for Sc-Zn (Table 7.4) 35 36 37 Lanthanides and Actinides All these elements have the configuration All [core] nsx (n - 1)dy (n - 2)fz and so are [core] (n f-block elements. f-block Cerium [Xe] 6s2 5d1 4f1 Uranium [Rn] 7s2 6d1 5f3 Lanthanide Element Lanthanide Configurations Configurations 4f orbitals used for Ce Lu and 5f for Th - Lr (Table 7.2) 38 39 40 Ion Configurations To form cations from elements remove 1 or To more e- from subshell of highest n [or highest (n + l)]. (n P [Ne] 3s2 3p3 - 3e- → P3+ [Ne] 3s2 3p0 3p 3p 3s 3s 2p 2p 2s 2s 1s 1s 41 Ion Configurations For transition metals, remove ns electrons and For then (n - 1) electrons. then Fe [Ar] 4s2 3d6 loses 2 electrons Fe2+ [Ar] 4s0 3d6 Fe2+ Fe 4s 3d 4s 3d Fe3+ To form cations, always remove electrons of highest n value first! 4s 3d 42 PERIODIC PERIODIC TRENDS TRENDS General Periodic Trends • Atomic and ionic size • Ionization energy • Electron affinity Larger orbitals. Electrons held less tightly. 43 44 Effective Nuclear Charge, Z* • Z* is the nuclear charge experienced by Z* the outermost electrons. See Figure 7.3 See the • Explains why E(2s) < E(2p) • Z* increases across a period owing to Z* incomplete shielding by inner electrons. incomplete • Estimate Z* = [ Z - (no. inner electrons) ] Estimate • Charge felt by 2s e- in Li Z* = 3 - 2 = 1 Charge • Be Z* = 4 - 2 = 2 Be Z* •B Z* = 5 - 2 = 3 and so on! Z* 45 Effective Effective Nuclear Charge Charge See Figure 7.3 Z* is the nuclear Z* charge experienced by the outermost electrons. electrons. Electron cloud for 1s electrons Effective Nuclear Charge, Z* • Atom • • • • • • • Li Be B C N O F Z* Experienced by Electrons in Valence Orbitals +1.28 ------+2.58 Increase in Z* Increase +3.22 across a +3.85 period period +4.49 +5.13 46 General Periodic Trends • Atomic and ionic size • Ionization energy • Electron affinity Larger orbitals. Electrons held less tightly. 47 48 Atomic Radii See Active Figure 7.8 49 Atomic Size • Size goes UP on going down a group. See Figure 7.8. See • Because electrons are Because added further from the nucleus, there is less attraction. attraction. • Size goes DOWN on going across a period. across 50 Atomic Size Size decreases across a period owing Size decreases to increase in Z*. Each added electron feels a greater and greater + charge. feels Large Small 51 Trends in Atomic Size See Active Figure 7.8 Radius (pm) 250 K 1st transition series 3rd period 200 Na 2nd period Li 150 Kr 100 Ar Ne 50 He 0 0 5 10 15 20 25 Atomic Number 30 35 40 Sizes of Transition Elements See Figure 7.9 52 Sizes of Transition Elements See Figure 7.9 • 3d subshell is inside the 4s 3d subshell. subshell. • 4s electrons feel a more or less 4s constant Z*. constant • Sizes stay about the same and Sizes chemistries are similar! chemistries 53 54 Ion Sizes Li,152 pm 3e and 3p Does the size go + up+or down when up , 60 pm Li l2e and 3 p electron osing an to form a cation? to 55 Ion Sizes + Li,152 pm 3e and 3p Li + , 78 pm 2e and 3 p Forming a Forming cation. cation. • CATIONS are SMALLER than the SMALLER atoms from which they come. atoms • The electron/proton attraction has The gone UP and so size DECREASES. DECREASES 56 Ion Sizes F,64 pm 9e and 9p Does the size go up or Does ­ down when gaining an electron F­ , 136 pm to form an anion? anion? 10 e and 9 p 57 Ion Sizes ­ F, 71 pm 9e and 9p F ­ , 133 pm 10 e and 9 p Forming an anion. an • ANIONS are LARGER than the atoms from LARGER which they come. which • The electron/proton attraction has gone The DOWN and so size INCREASES. INCREASES • Trends in ion sizes are the same as atom Trends sizes. 58 Trends in Ion Sizes See Active Figure 7.12 59 Redox Reactions Why do metals lose Why electrons in their reactions? Why does Mg form Mg2+ Why ions and not Mg3+? ions Why do nonmetals take Why on electrons? on 60 Ionization Energy IE = energy required to remove an electron IE from an atom in the gas phase. from Mg (g) + 738 kJ Mg+ (g) + eMg 61 Ionization Energy IE = energy required to remove an electron IE from an atom in the gas phase. from Mg (g) + 738 kJ → Mg+ (g) + eMg Mg+ (g) + 1451 kJ Mg2+ (g) + e- 62 Ionization Energy Ionization Mg (g) + 735 kJ → Mg+ (g) + eMg Mg+ (g) + 1451 kJ → Mg2+ (g) + e- Mg2+ (g) + 7733 kJ Mg3+ (g) + eEnergy cost is very high to dip into a shell Energy of lower n. Trends in Ionization Energy See Active Figure 7.10 63 64 Trends in Ionization Energy 1st Ionization energy (kJ/mol) 2500 He Ne 2000 Ar 1500 Kr 1000 500 0 1 H 3 Li 5 7 9 11 Na 13 15 17 19 K 21 23 25 27 29 31 Atomic Number 33 35 65 Trends in Ionization Energy • IE increases across a period IE because Z* increases. because • Metals lose electrons more Metals easily than nonmetals. easily • Metals are good reducing Metals agents. agents. • Nonmetals lose electrons with Nonmetals difficulty. difficulty. 66 Trends in Ionization Energy • IE decreases down a IE group • Because size increases. • Reducing ability generally Reducing increases down the periodic table. 67 Electron Affinity A few elements GAIN electrons few GAIN to form anions. anions Electron affinity is the energy Electron involved when an atom gains an electron to form an anion. an A(g) + e- A-(g) E.A. = ∆U 68 Trends in Electron Affinity See Active Figure 7.11 69 Trends in Electron Affinity • See Figure 7.11 and See Appendix F Appendix • Affinity for electron Affinity increases across a period (EA becomes more positive). more • Affinity decreases down Affinity a group (EA becomes Note effect of atom Note less positive). less size on F vs. Cl size ...
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This note was uploaded on 06/02/2011 for the course ACC 101 taught by Professor Duean during the Spring '11 term at Cornish.

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