Electrons in Atoms & Periodic Table

Transcription

1 Electrons in Atoms & Periodic Table 1

2 Where Are the Electrons? Quantum Theory or Wave Theory: a description of the electron configuration. One of the greatest achievements of mankind. 2

3 Arrangement of e - in Atoms Determines: chemical reactivity bonding between atoms Periodic Table many physical properties 3

4 Atomic Models: History Each atomic model was eventually replaced because of new experimental evidence

5 Dalton: 1803 Concept of the atom as smallest unit of an element. Indivisible particle 5

6 Thomson: 1897 Discovered the e - Atom has parts!! electron positive charge Plum pudding model 6

7 Rutherford: 1911 Au foil experiment Dense nucleus with positive charge Most of atom is empty space + Nuclear model 7

8 Nuclear Model: Problem Opposite charges attract, so what keeps the electrons and nucleus apart? + 8

9 Bohr: 1913 e - held in orbits Motion of e - keeps them from falling into nucleus Similar to planets around sun 9

10 Bohr: Planetary Model e - move in circular orbits around nucleus, and each orbit has a certain energy. 10

11 Bohr: Planetary Model E 3 E 2 E 1 + Quantized energy levels 11

12 Bright Line Spectrum of Hydrogen 12

13 energy Stair Analogy: H spectrum due to e - transition between orbits. E 5 E 4 Stairs are quantized. E 3 E 2 Not a ramp E 1 13

14 energy e - in Ground State E 5 E 4 Ground state is lowest energy of the e -. E 3 E 2 E 1 14

15 energy e - in Excited State E 5 e - absorbs energy to move to a higher energy level. E 4 E 3 E 2 E 1 15

16 energy e - in Excited State E 5 E 4 E 3 E 2 E 1 16

17 energy e - Returning to Ground E 5 e - gives off energy as light photon E 4 E 3 E 2 E light =E excited -E ground E 1 17

18 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. 18

19 energy e - Returning to Ground E 5 E 4 lower energy photon E 3 E 2 E 1 19

20 Bohr: Hydrogen Emission Spectrum E 3 E 2 E1 + e - absorbs energy (heat, elec.) e - falls to lower E and gives off energy as light E light =E 3 -E 1 20

21 Bohr Theory: Failings Why do e - only have certain orbit energies? Only explains the hydrogen atom exactly. 21

22 Quantum Mechanics (Wave Theory) 1926: E. Schrodinger Currently accepted theory Can not determine exact location of an electron! Wow! 22

23 Quantum Mechanics e - in atomic orbitals Can only determine the probability of locating an electron. electron cloud 23

24 Models + Dalton Thomson Rutherford + + Bohr Quantum 24

25 Atomic Orbital A region in space around the nucleus with high probability of finding an electron. Each atomic orbital can hold 2 e Analogy: student in a desk 25

26 Wave Model Each e - is arranged in an atom according to its energy. 3rd energy level (higher shell) + 2 nd energy level 1 st energy level (lowest shell) 26

27 Overview Bohr Quantum _embedded&v=8rohpz0a70i#t=4 27

28 Regents e - Notation Regents Periodic Table gives the number of electrons in each energy level or shell. 1 st shell - 2 nd shell - 3 rd shell [C] =

29 Regents e - Configuration What is the e - configuration for: 1. sodium 2. argon 3. calcium 4. copper 7. lead 6. radium (Ra) What is the maximum electrons in: Shell 1? Shell 2? Shell 3? 29

30 Regents e - Configuration What happens to the number of electrons in each shell going from: Ca to Sc? Zn to Ga? 30

31 Noble Gases At end of each row in Periodic Table are the noble or inert gases with 8 e - in the highest shell. Stable (not reactive) elements 31

32 Periodic Table by Shell Transition elements Inner transition elements 32

33 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 33

34 Valence Electrons Note all elements in a Group have same number of valence electrons. Group 1: Li, Na, K, Rb, Cs, Fr Group 16: O, S, Se, Te, Po This is why elements in the same group have similar properties. 34

35 Excited State Remember excited state? (e - have absorbed energy to move to a higher energy level) [Al] is ground state Al [Al]* could be excited state Al What element is ? 35

36 Flame Test (Lab) Adding energy can cause e - to jump from ground state (as written in Regents table) to excited state. When e - falls back, it emits light. 36

37 Flame Test for Copper Cu atom in excited state: Cu atom in ground state: Can return to ground state by emitting energy 37

38 Flame Test for Copper Which photon has greater energy: When an e - the falls from E 5 to E 3 or When an e - the falls from E 5 to E 2? 38

39 Emission Spectrum Flame test Neon signs Fireworks Fireplace colors 39

40 Bonding Electron configurations are the key to bonding. Atoms become ions to achieve Noble gas electron configuration. 40

41 Atoms vs. Ions F atom: [F] = 2-7 What does F need in order to have the e - configuration of a Noble gas? F - ion: [F - ] = 2-8 = [Ne] Na atom: [Na] = Na + ion: [Na + ] = 2-8 = [Ne] 41

42 Practice Write the e - configuration for Fe & Cu. Based on e - configuration, predict the charge of: Mg ion S ion Write two e - configurations for excited states of calcium. 42

43 Periodic Relationships 43

44 Early chemists describe the first element. 44

45 Tabulation of Elements Mendeleev (1869) Tabulated by chem. & physical properties Arranged by mass Predicted missing elements and properties 45

46 Modern Periodic Table Argon vs. potassium problem. Now ordered by atomic number, not mass. Element 101 (Md) 46

47 Periodic Table Most important tool in chemistry Key to understanding chemical and physical properties Each group has same electron configuration for outer shell. 47

48 Regents Periodic Table Elements arranged by atomic no. (#p + ) Symbol Atomic number & atomic mass Electron configuration Charges 48

49 Representative Elements Groups 1, 2 & Last digit of group number gives the number of valence electrons. Examples: oxygen Group 16 (6 VE) sodium Group 1 (1 VE) 49

50 Representative Elements Some groups have special names Group 1: alkali metals Group 2: alkaline earth metals Group 17: halogens 50

51 Noble Gases Last element in each Period 8 VE, except He very stable (non-reactive) 51

52 Transition Elements e.g. Iron Regents: e - filling Compounds with these elements have colored solutions. 52

53 Inner Transition Elements At bottom of Periodic Table for convenience. 53

54 Trends in Atomic Size Atomic size is measured by radius. Table S R For chlorine: radius = 100. pm What is its radius in meters? 54

55 Atomic Radius: Trends????????? OK (model) 55

56 Atomic Radius Down a Group: size increases due to adding electrons to higher energy levels (shells) further from the nucleus. 56

57 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 + =

58 Atomic Size: Across a Period smaller Boron (2-3) vs. Carbon (2-4) 58

59 larger Atomic Radius smaller Row: greater nuclear charge Column: e - in higher shell 59

60 Atomic Radius Try It: Arrange these atoms in order of increasing size. N, O, P, S O < N < S < P 60

61 Ionization Energy (I) Chemical properties determined by valence electrons. First ionization energy (I): energy (kj/mol) to remove a valence e - from an atom. If ionization energy is high, e - held tightly. 61

62 Ionization Energy I is endothermic (need to put energy in to pull off an e - ) I + X(g) X + (g) + e - ionization energy 62

63 Ionization Energy: Table S I 1 I 1 across Period down Group I 1 Atomic Number 63

64 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? 64

65 I Predicts Ionic Charges Element I 1 (kj/mol) I 2 (kj/mol) I 3 (kj/mol) Na Mg Na atom lose 1 e- Na + ion 2-8 Mg atom lose 2 e- Mg +2 ion

66 Ionization Energy Which has smaller I and why? O or S Ge or Br 66

67 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 67

68 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 68

69 (model) Ionic Radii cations anions 69

70 Ionic Radii Place in order of increasing size. Fe, Fe 2+ and Fe 3+ 70

71 Try It!!! 1. Use e - configuration to predict the charge of Ca ion. 2. Is this ion larger or smaller than its atom? 71

72 Electronegativity The tendency of an atom to attract bonding electrons. Water: which atom wins the battle for the bonding e -? H O H 72

73 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 73

74 Electronegativity Why don t the Noble gases have electronegativity values? 74

75 Electronegativity Example: Water 3.4 O d- O slightly H H H H d+ d+ Water is a polar molecule. 75

76 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. (Same trend as ionization energy.) 76

77 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) 77

78 Diagonal Relationships Smallest R Largest I 1 Largest EN Least metallic Largest R Smallest I 1 Smallest EN Most metallic 78

79 79

80 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? 80

81 Warm-up What was Bohr s explanation for the emission or bright-line spectrum of hydrogen? + 81

82 Warm-up What is the name of the region outside the nucleus where electrons are most probably found? 82

83 Warm-up What is the name of the region outside the nucleus where electrons are most probably found? Write the Regents electron configuration for arsenic. What does each of the numbers mean? How many valence electrons does arsenic have? 83

84 Warm-up What is 2-8-7? What is 2-7-8? What is the e - configuration of gold? How many valence electrons does manganese have? What is the electron configuration of the nitride ion? 84

85 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? -across a row? Why? What is e - config. of Al +3? 85

86 Warm-up Which element, P or S, is bigger (larger radius)? Explain. 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 86

87 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? 87

88 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 88

89 Element Song /elements.html 89