Electrons in Atoms. Quantum Theory or Wave Theory. It s Unreal!! Check your intuition at the door. description of the electronic structure in atoms
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1 Electrons in Atoms It s Unreal!! Check your intuition at the door. Quantum Theory or Wave Theory description of the electronic structure in atoms 1
2 Quantum Theory Unlike anything in our macroscopic world. wave-particle duality probability of electron location orbital shapes 2
3 Quantum Theory 1900~ 1930 One of the greatest achievements of mankind. 3
4 e - Arrangement in Atoms Determines: chemical reactivity bonding between atoms Periodic Table many physical properties 4
5 Atomic Models: History Each atomic model was eventually replaced in light of new experimental evidence
6 Dalton: 1803 Concept of the atom as smallest unit of an element. Indivisible particle 6
7 Thomson: 1897 Discovered the e - Atom has parts!! electron +charge Plum pudding model 7
8 Rutherford: 1911 Au foil experiment Nucleus with positive charge Most of atom is empty space + Nuclear model 8
9 Nuclear Model: Problem + What keeps the electrons and nucleus apart? 9
10 Neils Bohr: 1913 e - held in orbits Motion of e - keeps them from falling into nucleus Similar to planets around sun 10
11 Bohr: Planetary Model e - move in circular orbits around nucleus, and each orbit has a certain energy. 11
12 Bohr: Planetary Model + E 3 E 2 E 1 Quantized energy levels 12
13 Bright Line Spectrum of Hydrogen 13
14 Stair Analogy: H spectrum due to e - transitions. energy E5 E 4 E 3 Stairs are quantized. E 2 Not a ramp E 1 14
15 energy e - in Ground State E 5 E 4 Ground state is lowest energy of the e -. E 3 E 2 E 1 15
16 e - in Excited State energy E5 E 4 E 3 e - absorbs energy to move to a higher energy level. E 2 E 1 16
17 e - in Excited State energy E5 E 4 E 3 E 2 E 1 17
18 e - Returning to Ground energy E5 E 4 e - gives off energy as light E 3 E 2 photon E light =E excited -E ground E 1 18
19 E light =E excited -E ground The energy of the light is the difference between the higher and lower energy level of the electron. Each energy of light corresponds to a unique color of light. 19
20 energy e - Returning to Ground E 5 E 4 lower energy photon E 3 E 2 E 1 20
21 Bohr: Hydrogen Emission Spectrum E 3 E 2 E 1 + e - absorbs energy (heat, elec.) e - falls to lower E and gives off energy as light E light =E 3 -E 1 21
22 Emission Spectrum Flame test Neon signs Fireworks Fireplace colors 22
23 Bohr Theory: Failings Why do e - only have certain orbital energies? Only explains the hydrogen atom exactly. 23
24 Quantum Mechanics Or Wave Model 1926: E. Schrodinger e - location (orbital) described by a probability function: Y 24
25 Schrodinger Equation d 2 Y dx 2 + d2 Y dy 2 d 2 Y + + dz 2 8p 2 m h 2 (E U) Y = 0 Many solutions. 25
26 Quantum Mechanical Model or Wave Model e - in atomic orbitals (2e - ) Can only determine the probability of locating an electron e - cloud 26
27 Models + Dalton Thomson Rutherford Bohr Quantum 27
28 Atomic Orbital A region in space around the nucleus with high probability of finding an electron. Each atomic orbital can hold 2 e - maximum - - Analogy: student in a desk 28
29 e - orbital (location) determined by: 1. Principal quantum number (shell) (principle energy level) 2. Sublevel (subshell) Only certain combinations are allowed. 29
30 Principal Quantum Number (n) or Principal Energy Level n = 1, 2, 3, 4, 5 (integer values) Gives overall energy of an e - and its distance from nucleus. 30
31 Energy Sublevels (subshell) Gives shape of the e - cloud s, p, d, f 31
32 Energy Sublevels s has 1 orbital What is the p has 3 orbitals max. number of e - in a d has 5 orbitals d orbital? f has 7 orbitals Each orbital can hold 2 e - 32
33 Sublevel Shapes s spherical 90% p x p y p z dumb bell 33
34 Allowed Combinations n Orbitals # Orbitals #e 1 1s s 2p s 3p 3d s 4p 4d 4f (model) 34
35 e - Configurations: 3 Rules 1. Aufbau: arrange e - by lowest energy level first 2. Pauli Exclusion Principle: only 2 e - per orbital 3. Hund s Rule: maximize the number of parallel spin e - when filling a sublevel 35
36 1. Aufbau: Arrange e - by lowest energy level first Increasing Energy 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s lower than 3d!! 36
37 3d hold 10e - 4s hold 2e - 3p hold 6e - 3s hold 2e - 2p hold 6e - 2s hold 2e - 1s hold 2e - 37
38 Orbital Box Diagram energy 1s 2s 2p 3s 3p 4s energy continued 3d 4p 5s 38
39 2. Pauli Exclusion Principle: only 2 e - per orbital Electrons have spin! 1s orbital Paired electrons: e - in an orbital have opposite spins 39
40 3. Hund s Rule Maximize parallel (same) spins when filling a sublevel In a sublevel, put one e - in each orbital before pairing Example: 4 e - in a p sublevel 40
41 Overview edded&v=8rohpz0a70i#t=4 41
42 Hydrogen: 1 electron 1s 2s 2p 3s Recall for neutral atom #e - is same as # p + (atomic number). 42
43 Electron Configuration Hydrogen 1s # of e - s 1s 1 n sublevel 43
44 Regents Table Notation Regents Periodic Table gives only the number of electrons in each principle energy level, n. 1 st level - 2 nd level - 3 rd level Hydrogen 1s 1 Regents 1 44
45 Helium: 2 electrons 1s 2s 2p 3s Notation: Regents 1s
46 Lithium: 3 electrons 1s 2s 2p 3s Notation: Regents 1s 2 2s
47 Boron: 5 electrons 1s 2s 2p 3s Notation: Regents 1s 2 2s 2 2p
48 Nitrogen: 7 electrons 1s 2s 2p 3s Notation: Regents 1s 2 2s 2 2p
49 Fluorine: 9 electrons 1s 2s 2p 3s Notation: Regents 1s 2 2s 2 2p
50 Neon: 10 electrons 1s 2s 2p 3s Notation: Regents 1s 2 2s 2 2p
51 Sodium: 11 electrons 1s 2s 2p 3s Notation: Regents 1s 2 2s 2 2p 6 3s
52 Shorthand for Sodium 1s 2 2s 2 2p 6 3s 1 [Ne] 3s 1 Use the preceding [noble gas] 52
53 Other 3 rd Period Elements energy 1s 2s 2p 3s 3p 4s Al 1s 2 2s 2 2p 6 3s 2 3p Ar 1s 2 2s 2 2p 6 3s 2 3p
54 Recall: 4s Out-of-Order Energy 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s lower than 3d!! So the next element, K, is rather than
55 Other Elements K [Ar]4s Sc [Ar]4s 2 3d Cr [Ar]4s 1 3d weird 55
56 Noble Gases At end of each row in Periodic Table are the noble or inert gases with filled ns and np orbitals. Stable (not reactive) elements 56
57 Periodic Table by Subshell 1s 2s Transition elements filling d orbitals 2p 1s 3s 3p 4s 3d 4p 5s 4d 5p 6s 5d 6p 7s 6d 4f 5f Inner transition elements filling f orbitals 57
58 Let s Predict One Using the Periodic Table, predict the full, shorthand, and Regents e - notation for zinc. 58
59 You Try It!! 1. Predict the full, shorthand, & Regents e - configuration of: Ca As 2.How many unpaired electrons are there in these elements? 3. Write the shorthand notation for: Zr 59
60 Valence Electrons Electrons that are in the highest energy level are called valence electrons. These are the most important electrons when atoms bond. Why? How many valence electrons in: Li Fe Cu 60
61 Excited State Remember excited state? (e - have absorbed energy to move to a higher energy level) What atom is 1s 2 2s 2 2p 3 3s 1? What atom is ? 61
62 Flame Test for Copper Cu atom in excited state: Cu atom in ground state: Can return to ground state by emitting energy as light 62
63 Weirdos! Some ground state elements on the Periodic Table have 1 or 2 electrons out of order. e.g. Cu and Cr You are not responsible for these. 63
64 Bonding (the octet rule) The e - configurations are the key to bonding. Some atoms will become ions to achieve Noble gas electron configuration. 64
65 F atom vs. F - ion 1s 2s 2p F F - extra e - = [Ne] 65
66 Na atom vs. Na + ion Na 1s 2s 2p 3s Na + missing e - = [Ne] 66
67 Practice Write full and Regents electron configuration for: Mn, Cu, Tc, S -2 67
68 Periodic Relationships 68
69 Early chemists describe the first element. 69
70 Tabulation of Elements Mendeleev (1869) Tabulated by chem. & physical properties Arranged by mass Predicted missing elements and properties 70
71 Modern Periodic Table Argon vs. potassium problem. Now ordered by atomic number, not mass. Element 101 (Md) 71
72 Periodic Table Most important tool in chemistry Key to understanding chemical and physical properties Electron configurations: key to grouping the elements 72
73 Representative Elements Groups 1, 2 & Filling the s or p subshells Last digit of group number gives the number of valence electrons. 73
74 Representative Elements Some groups have special names Group 1: alkali metals Group 2: alkaline earth metals Group 17: halogens 74
75 Noble Gases filled p subshells (except He) ns 2 np 6 (8 valence e - ) e.g. neon 1s 2 2s 2 2p 6 very stable (non-reactive) 75
76 Transition Elements filling d subshell d block e.g. Iron [Ar]4s 2 3d 6 Regents: Salts yield colored solutions. 76
77 Inner Transition Elements filling the f subshell f block 77
78 s 2s 3s Periodic Table by Subshell 4s 3d 4p 5s 4d 5p 6s 5d 6p 7s 6d 4f 1s 2p 3p 5f 78
79 Trends in Atomic Size Atomic size is measured by radius. Table S R For chlorine: = 100. pm =? m 79
80 Atomic Radius: Trends????????? OK (model) 80
81 Atomic Radius Down a Group: size increases due to adding electrons to higher energy levels (shells) further from the nucleus. 81
82 Atomic Size: Across a Period Electrons added to same shell Nuclear charge increases (more p + ) Greater inward pull on the electrons Atoms get smaller p + =
83 Atomic Size: Across a Period smaller Boron (2-3) vs. Carbon (2-4) 83
84 larger Atomic Radius smaller Row: greater nuclear charge Column: e - in higher shell 84
85 Atomic Radius Try It: Arrange these atoms in order of increasing size. N, O, P, S O < N < S < P 85
86 Ionization Energy (I) Chemical properties determined by valence electrons. Ionization energy: energy (kj/mol) to remove an e - from an atom. If I high, e - held tightly. 86
87 Ionization Energy I is endothermic (need to put energy in to pull off an e - ) 1 st ionization energy or I 1 energy + X(g) X + (g) + e - 87
88 Ionization Energy: Table S I 1 I 1 across Period down Group I 1 Atomic Number 88
89 Trends in I (due to size) I 1 decreases going down a Group. The e - are farther from the nucleus. I 1 increases going across a Period. The e - are closer to the nucleus. Which corner of Periodic Table has: -highest I 1? -lowest I 1? 89
90 I Predicts Ionic Charges Element I 1 (kj/mol) I 2 (kj/mol) I 3 (kj/mol) Na Mg Na 1s 2s 2p 3s 90
91 Ionization Energy Which has smaller I 1 and why? O or S P or Cl 91
92 Trends in Ionic Size Cation is smaller than its atom. (less e - with same # protons) Na -1e - Na pm 95 pm Al -3e - Al pm 50 pm 92
93 Trends in Ionic Size Anion is larger than its atom. (more e - with same # protons) Cl Cl +1e pm 181 pm F +1e - F - 60 pm 136 pm 93
94 (model) Ionic Radii cations anions 94
95 Ionic Radii Place in order of increasing size. Fe, Fe 2+ and Fe 3+ 95
96 Try It!!! 1.Use e - configuration to predict the charge of Ca ion. 2.Is this ion larger or smaller than its atom? 96
97 Electronegativity The tendency of an atom to attract bonding electrons. Water: which atom wins the battle for the bonding e -? H O H 97
98 Electronegativity An arbitrary scale from 0 to Least EN Most EN Fr (0.7) F (4.0) Low attraction for e - in bond High attraction for e - in bond 98
99 Electronegativity Why don t the Noble gases have electronegativity values? 99
100 Electronegativity Example: Water 3.4 O d- O slightly H H H H d+ d+ Water is a polar molecule. 100
101 Electronegativity Group Trend: EN decreases going down a group. Atoms get larger, so bonding e - are farther from the nucleus. Period Trend: EN increases going across a period. Atoms get smaller, so bonding e - are closer to the nucleus. (Similar to ionization energy.) 101
102 Metallic Character Metals lose e - to become cations. Which element is the most metallic? (smallest ionization energy) Nonmetals gain e - to become anions. Which element is the least metallic? (largest ionization energy) 102
103 Diagonal Relationships Smallest R Largest I 1 Largest EN Least metallic Largest R Smallest I 1 Smallest EN Most metallic 103
104 104
105 Warm-up What did Rutherford s gold foil experiment show about the structure of the atom? How did Bohr s model of the atom differ from the prior model of the atom? 105
106 Warm-up What was Bohr s explanation for the emission or bright-line spectrum of hydrogen? + 106
107 Warm-up What two quantum properties determine the location of an electron in an atom? electron neutron proton 107
108 Warm-up What are the 3 rules for placing e- around an atom? What is the order of energy sublevels? Complete the chart: n Sublevels No. orbitals No. e What is the relationship between the value of n and the number of e-? 108
109 Warm-up For oxygen & sulfur write: box diagram electron configuration shorthand notation Regents configuration 109
110 Warm-up Write the shorthand e - config. for: Iron (26) Rhodium (45) What is the max. number of e - that can be on the 4 rd principal quantum number in the ground state? 110
111 Warm-up What are the names of Groups: 1, 2, 17, and 18. How many valence e - in Co? What is the trend size: -down a group? Why? -across a row? What is e - config. of Al +3? 111
112 Warm-up Define first ionization energy, I 1. What is the trend in I 1 across a row and down a group? Explain. Place the following elements in order of increasing I 1 : P, Cl, As 112
113 Warm-up What is metallic character? How is metallic character related to ionization energy? What happens to metallic character going down Group 15? Which has greater metallic character: Fe or Na? 113
114 Warm-up Define each term, state the trend, and explain why: Atomic radius across a row Ionization Energy down a group Electronegativity across a row Metallic character down a group 114
115 Element Song /elements.html 115
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