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Chemistry and Chemical Reactivity 6th Edition 1 John C. Kotz Paul M. Treichel Gabriela C. Weaver CHAPTER 8 Atomic Electron Configurations and Chemical Periodicity Lectures written by John Kotz ©2006 2006 Brooks/Cole Thomson © Brooks/Cole - Thomson ATOMIC ELECTRON CONFIGURATIONS AND PERIODICITY © 2006 Brooks/Cole - Thomson 2 Arrangement of Electrons in Atoms Electrons in atoms are arranged as SHELLS (n) SUBSHELLS (l) ORBITALS (ml) © 2006 Brooks/Cole - Thomson 3 9 Pauli Exclusion Principle No two electrons in the same atom can have the same set of 4 quantum numbers. That is, each electron has a unique address. © 2006 Brooks/Cole - Thomson Assigning Electrons to Subshells • In H atom all subshells of same n have same energy. • In many-electron atom: a) subshells increase in energy as value of increases. S<p<d<f b) for shells increase in energy: 1<2<3<4<… © 2006 Brooks/Cole - Thomson 15 16 Electron Filling Order Figure 8.5 © 2006 Brooks/Cole - Thomson 19 Writing Atomic Electron Configurations Two ways of writing configs. One is called the spdf notation. spdf notation for H, atomic number = 1 1 1s value of n © 2006 Brooks/Cole - Thomson no. of electrons value of l Writing Atomic Electron Configurations Two ways of writing configs. Other is called the orbital box 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 Other electron has n = 1, l = 0, ml = 0, ms = - 1/2 © 2006 Brooks/Cole - Thomson 20 21 See “Toolbox” on CD for Electron Configuration tool. © 2006 Brooks/Cole - Thomson 22 Electron Configurations and the Periodic Table Active Figure 8.7 © 2006 Brooks/Cole - Thomson 23 The electron configuration for chlorine is: 1. 2. 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 43 44 45 46 47 49 50 51 52 53 54 55 56 57 58 59 60 63 64 65 66 4. 67 48 61 3.4262 68 69 70 71 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86 87 88 89 90 91 92 93 94 95 96 97 98 99 100 101 102 103 104 105 106 107 108 109 110 111 112 113 114 115 116 117 118 119 120 121 122 123 124 125 126 127 128 129 130 131 132 133 134 135 136 137 138 139 140 141 142 143 144 145 146 147 148 149 150 151 152 153 154 155 156 157 158 159 160 161 162 163 164 165 166 167 168 169 170 171 172 0% 0% 0% 0% 173 174 175 176 177 178 179 181 182 183 184 185 186 187 188 189 190 191 192 193 194 195 196 197 198 199 200 201 202 203 204 205 206 207 208 209 210 211 212 213 214 215 216 217 218 219 220 221 222 223 224 225 226 227 228 229 230 © 2006 Brooks/Cole - Thomson 1 2 3 4 180 Lithium Group 1A Atomic number = 3 1s22s1 ---> 3 total electrons 3p 3s 2p 2s 1s © 2006 Brooks/Cole - Thomson 24 25 Beryllium 3p 3s 2p 2s 1s © 2006 Brooks/Cole - Thomson Group 2A Atomic number = 4 1s22s2 ---> 4 total electrons Boron 3p 3s 2p 2s 1s © 2006 Brooks/Cole - Thomson Group 3A Atomic number = 5 1s2 2s2 2p1 ---> 5 total electrons 26 Carbon 3p 3s 2p 2s 1s © 2006 Brooks/Cole - Thomson Group 4A Atomic number = 6 1s2 2s2 2p2 ---> 6 total electrons Here we see for the first time HUND’S RULE. When placing electrons in a set of orbitals having the same energy, we place them singly as long as possible. 27 28 Nitrogen 3p 3s 2p 2s 1s © 2006 Brooks/Cole - Thomson Group 5A Atomic number = 7 1s2 2s2 2p3 ---> 7 total electrons 29 Oxygen 3p 3s 2p 2s 1s © 2006 Brooks/Cole - Thomson Group 6A Atomic number = 8 1s2 2s2 2p4 ---> 8 total electrons 30 Fluorine Group 7A Atomic number = 9 1s2 2s2 2p5 ---> 9 total electrons 3p 3s 2p 2s 1s © 2006 Brooks/Cole - Thomson 31 Neon Group 8A Atomic number = 10 1s2 2s2 2p6 ---> 10 total electrons 3p 3s 2p 2s 1s © 2006 Brooks/Cole - Thomson Note that we have reached the end of the 2nd period, and the 2nd shell is full! 32 1. 2. 3. 4. W Mo Ru Pm 0% 1 © 2006 Brooks/Cole - Thomson 0% 0% 2 3 0% 4 33 Electron Configurations of p-Block Elements © 2006 Brooks/Cole - Thomson Sodium 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. All Group 1A elements have [core]ns1 configurations. © 2006 Brooks/Cole - Thomson 34 35 Aluminum Group 3A Atomic number = 13 1s2 2s2 2p6 3s2 3p1 [Ne] 3s2 3p1 All Group 3A elements have [core] ns2 np1 configurations where n is the period number. 3p 3s 2p 2s 1s © 2006 Brooks/Cole - Thomson 36 Phosphorus Yellow P Group 5A Atomic number = 15 1s2 2s2 2p6 3s2 3p3 [Ne] 3s2 3p3 All Group 5A elements have [core ] ns2 np3 configurations where n is the period number. Red P 3p 3s 2p 2s 1s © 2006 Brooks/Cole - Thomson 37 Calcium Group 2A Atomic number = 20 1s2 2s2 2p6 3s2 3p6 4s2 [Ar] 4s2 All Group 2A elements have [core]ns2 configurations where n is the period number. © 2006 Brooks/Cole - Thomson Electron Configurations and the Periodic Table © 2006 Brooks/Cole - Thomson 38 39 1. 2. 3. 4. Hf Lu Pb Sn 0% 1 © 2006 Brooks/Cole - Thomson 0% 0% 2 3 0% 4 Transition Metals Table 8.4 All 4th period elements have the configuration [argon] nsx (n - 1)dy and so are d-block elements. Chromium © 2006 Brooks/Cole - Thomson Iron Copper 40 Transition Element Configurations 3d orbitals used for Sc-Zn (Table 8.4) © 2006 Brooks/Cole - Thomson 41 42 © 2006 Brooks/Cole - Thomson 43 1. 2. 3. 4. Zn Ca Ge Ni 0% 1 © 2006 Brooks/Cole - Thomson 0% 0% 2 3 0% 4 44 The electron configuration for tin, Sn, element 50, is: 1. 2. 3. 4. 0% 1 © 2006 Brooks/Cole - Thomson 0% 0% 2 3 0% 4 Lanthanides 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 © 2006 Brooks/Cole - Thomson 45 Lanthanide Element Configurations 4f orbitals used for Ce - Lu and 5f for Th - Lr (Table 8.2) © 2006 Brooks/Cole - Thomson 46 47 © 2006 Brooks/Cole - Thomson 52 PERIODIC TRENDS © 2006 Brooks/Cole - Thomson General Periodic Trends • Atomic and ionic size • Ionization energy • Electron affinity Higher effective nuclear charge Electrons held more tightly Larger orbitals. Electrons held less tightly. © 2006 Brooks/Cole - Thomson 53 Effective Nuclear Charge, Z* • Z* is the nuclear charge experienced by the outermost electrons. See Figure 8.6 and and Screen 8.6. • Explains why E(2s) < E(2p) • Z* increases across a period owing to incomplete shielding by inner electrons. • Estimate Z* by --> [ Z - (no. inner electrons) ] • Charge felt by 2s e- in Li Z* = 3 - 2 = 1 • Be Z* = 4 - 2 = 2 • B Z* = 5 - 2 = 3 and so on! © 2006 Brooks/Cole - Thomson 54 55 Effective Nuclear Charge Figure 8.6 Z* is the nuclear charge experienced by the outermost electrons. Electron cloud for 1s electrons © 2006 Brooks/Cole - Thomson 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 +3.22 Z* across a +3.85 period +4.49 +5.13 [Values calculated using Slater’s Rules] © 2006 Brooks/Cole - Thomson 57 General Periodic Trends • Atomic and ionic size • Ionization energy • Electron affinity Higher effective nuclear charge Electrons held more tightly Larger orbitals. Electrons held less tightly. © 2006 Brooks/Cole - Thomson 59 60 Atomic Radii Active Figure 8.11 © 2006 Brooks/Cole - Thomson 61 Atomic Size • Size goes UP on going down a group. See Figure 8.9. • Because electrons are added further from the nucleus, there is less attraction. • Size goes DOWN on going across a period. © 2006 Brooks/Cole - Thomson 62 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* © 2006 Brooks/Cole - Thomson 63 Trends in Atomic Size See Active Figure 8.11 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 © 2006 Brooks/Cole - Thomson 30 35 40 Sizes of Transition Elements See Figure 8.12 © 2006 Brooks/Cole - Thomson 64 Sizes of Transition Elements See Figure 8.12 • 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! © 2006 Brooks/Cole - Thomson 65 Density of Transition Metals 25 20 6th period Density (g/mL) 15 10 5th period 4th period 5 0 3B 4B 5B 6B 7B 8B Group © 2006 Brooks/Cole - Thomson 1B 2B 66 Ion Sizes Li,152 pm 3e and 3p © 2006 Brooks/Cole - Thomson Does+ the size go up+ or down Li , 60 pm when an 2e and 3losing p electron to form a cation? 67 68 Ion Sizes + Li,152 pm 3e and 3p Li + , 78 pm 2e and 3 p Forming a cation. • CATIONS are SMALLER than the atoms from which they come. • The electron/proton attraction has gone UP and so size DECREASES. © 2006 Brooks/Cole - Thomson Ion Sizes Does the size go up or down when gaining an electron to form an anion? © 2006 Brooks/Cole - Thomson 69 70 Ion Sizes F, 71 pm 9e and 9p F- , 133 pm 10 e and 9 p Forming an anion. • ANIONS are LARGER than the atoms from which they come. • The electron/proton attraction has gone DOWN and so size INCREASES. • Trends in ion sizes are the same as atom sizes. © 2006 Brooks/Cole - Thomson Trends in Ion Sizes Active Figure 8.15 © 2006 Brooks/Cole - Thomson 71 72 Compare the elements Na, B, Al, and C with regard to the following properties: Which has the largest atomic radius? 1. 2. 3. 4. Na B Al C 0% 1 © 2006 Brooks/Cole - Thomson 0% 0% 2 3 0% 4 73 Compare the elements Na, B, Al, and C with regard to the following properties: Which has the largest (most negative) electron affinity? 1. 2. 3. 4. Na B Al C 0% 1 © 2006 Brooks/Cole - Thomson 0% 0% 2 3 0% 4 74 Compare the elements Na, B, Al, and C with regard to the following properties: Which has the largest (most positive) ionization energy? 1. 2. 3. 4. Na B Al C 0% 1 © 2006 Brooks/Cole - Thomson 0% 0% 2 3 0% 4 75 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? © 2006 Brooks/Cole - Thomson Ionization Energy See CD Screen 8.12 IE = energy required to remove an electron from an atom in the gas phase. Mg (g) + 738 kJ ---> Mg+ (g) + e- © 2006 Brooks/Cole - Thomson 76 Ionization Energy See Screen 8.12 IE = energy required to remove an electron from an atom in the gas phase. Mg (g) + 738 kJ ---> Mg+ (g) + e- Mg+ (g) + 1451 kJ ---> Mg2+ (g) + eMg+ has 12 protons and only 11 electrons. Therefore, IE for Mg+ > Mg. © 2006 Brooks/Cole - Thomson 77 Ionization Energy See Screen 8.12 Mg (g) + 735 kJ ---> Mg+ (g) + eMg+ (g) + 1451 kJ ---> Mg2+ (g) + e- Mg2+ (g) + 7733 kJ ---> Mg3+ (g) + eEnergy cost is very high to dip into a shell of lower n. This is why ox. no. = Group no. © 2006 Brooks/Cole - Thomson 78 Trends in Ionization Energy Active Figure 8.13 © 2006 Brooks/Cole - Thomson 80 81 Trends in Ionization Energy 1st Ionization energy (kJ/mol) 2500 He Ne 2000 Ar 1500 Kr 1000 500 0 1 H 3 Li © 2006 Brooks/Cole - Thomson 5 7 9 11 Na 13 15 17 19 K 21 23 25 27 29 31 Atomic Number 33 35 Orbital Energies As Z* increases, orbital energies “drop” and IE increases. CD-ROM Screens 8.9 - 8.13, Simulations © 2006 Brooks/Cole - Thomson 82 Trends in Ionization Energy • 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. © 2006 Brooks/Cole - Thomson 83 84 Trends in Ionization Energy • IE decreases down a group • Because size increases. • Reducing ability generally increases down the periodic table. • See reactions of Li, Na, K © 2006 Brooks/Cole - Thomson 85 Which of the following groups of elements is arranged correctly in order of increasing first ionization energy? 1. 2. 3. 4. Mg < C < N < F N < Mg < C < F Mg < N < C < F F < C < Mg < N 0% 1 © 2006 Brooks/Cole - Thomson 0% 0% 2 3 0% 4 Which of the following elements would have the greatest difference between the first and the second ionization energy? 1. 2. 3. 4. lithium carbon fluorine nitrogen 0% 1 © 2006 Brooks/Cole - Thomson 0% 0% 2 3 0% 4 86 87 Which of the following groups of elements is arranged correctly in order of increasing first ionization energy? 1. 2. 3. 4. B < O < Al < F B < O < F < Al Al < B < O < F F < O < B < Al 0% 1 © 2006 Brooks/Cole - Thomson 0% 0% 2 3 0% 4 88 Which of the following groups of elements is arranged correctly in order of increasing affinity for electrons (that is, electron affinity becomes more negative)? 1. 2. 3. 4. Mg < S < Al < Cl Mg < Al < S < Cl Al < Mg < S < Cl Cl < S < Mg < Al 0% 1 © 2006 Brooks/Cole - Thomson 0% 0% 2 3 0% 4 89 Which of the following groups of elements is arranged correctly in order of decreasing atomic radius? 1. 2. 3. 4. Mg > S > Al > Cl Mg > Al > S > Cl Al > Mg > S > Cl Cl > S > Mg > Al 0% 1 © 2006 Brooks/Cole - Thomson 0% 0% 2 3 0% 4 91 Electron Affinity A few elements GAIN electrons to form anions. Electron affinity is the energy involved when an atom gains an electron to form an anion. A(g) + e- ---> A-(g) E.A. = ∆E © 2006 Brooks/Cole - Thomson 92 Electron Affinity of Oxygen O atom [He] + electron O- ion [He] EA = - 141 kJ © 2006 Brooks/Cole - Thomson ∆E is EXOthermic because O has an affinity for an e-. 93 Electron Affinity of Nitrogen N atom [He] + electron N- ion [He] EA = 0 kJ © 2006 Brooks/Cole - Thomson ∆E is zero for Ndue to electronelectron repulsions. Trends in Electron Affinity Active Figure 8.14 © 2006 Brooks/Cole - Thomson 94 95 Trends in Electron Affinity • See Figure 8.14 and Appendix F • Affinity for electron increases across a period (EA becomes more positive). • Affinity decreases down a group (EA becomes less positive). © 2006 Brooks/Cole - Thomson Atom EA F +328 kJ Cl +349 kJ Br +325 kJ I +295 kJ Note effect of atom size on F vs. Cl 96 Which of the following groups of elements is arranged correctly in order of increasing first ionization energy? 1. 2. 3. 4. Mg < C < N < F N < Mg < C < F Mg < N < C < F F < C < Mg < N 0% 1 © 2006 Brooks/Cole - Thomson 0% 0% 2 3 0% 4 Which of the following elements would have the greatest difference between the first and the second ionization energy? 1. 2. 3. 4. lithium carbon fluorine nitrogen 0% 1 © 2006 Brooks/Cole - Thomson 0% 0% 2 3 0% 4 97 98 Which of the following groups of elements is arranged correctly in order of increasing first ionization energy? 1. 2. 3. 4. B < O < Al < F B < O < F < Al Al < B < O < F F < O < B < Al 0% 1 © 2006 Brooks/Cole - Thomson 0% 0% 2 3 0% 4 99 Which of the following groups of elements is arranged correctly in order of increasing affinity for electrons (that is, electron affinity becomes more negative)? 1. 2. 3. 4. Mg < S < Al < Cl Mg < Al < S < Cl Al < Mg < S < Cl Cl < S < Mg < Al 0% 1 © 2006 Brooks/Cole - Thomson 0% 0% 2 3 0% 4 100 Which of the following groups of elements is arranged correctly in order of decreasing atomic radius? 1. 2. 3. 4. Mg > S > Al > Cl Mg > Al > S > Cl Al > Mg > S > Cl Cl > S > Mg > Al 0% 1 © 2006 Brooks/Cole - Thomson 0% 0% 2 3 0% 4