# Chapters 4,5,6

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• Last pointwill be Illustrated with Bead Activity later in the week.
• In #3, they can be rearranged also includes combined and separated.
• I will show “The Atom – Episode 1” in an upcoming class which shows the history of the atom. Students are responsible for answering the questions and knowing it for the test.
• What the above relative masses show is that an electron is ~1820 times smaller (by mass) than an proton or neutron. Use the money equivalent to explain. A proton/neutron is worth \$20 and an electron is worth \$0.01.
• After this slide have students do Atomic Number/Mass Number activity for CW/HW.In Chromium row, #p is 1st, #e is 2nd, #n is thirdIn 2nd row (As) row, the order is Atomic # (33), #electrons (33), Mass #(75), As-75 IsotopeIn 3rd row (Rn-222), the order is # protons (86), atomic #(86), Isotope (Rn-222), #N°
• QPA means Quality Point Average.
• Note: An 85 in chemistry would give 510 Quality points and raise your average to 90.90. Because a difference of 10 in chem means a little bit more than and difference in 10 of English. 60 QP vs 50 QPs.
• Going back to our QPA analogy, the mass is like the grade and the abundance is like the credit. The mass x abundance is like the QP points. When you divide the Quality Points (grade*credit) by credit you get a weighed grade. When you divide the sum of the Mass x abundance numbers by total abundance (which should always be 100) you get the weighted atomic mass.
• The tabular format in the previous slide shows the method assuming 100 atoms. This is a more direct method, first show changing the percents to a fraction and then multiplying by mass. You are just going to get the mass of 1 atom in this way.Either way is acceptable. It’s really just a matter of where you are dividing by 100.
• Other Atomic Mass abundances:Boron-10 : 19.8% Mass: 10.01 Boron-11 : 80.2% 11.01Silicon-28 : 92.23% Mass: 27.98Silicon-29 : 4.67% 28.98Silicon-30 : 3.10% 29.98Iron-54 : 5.8% Mass: 53.94Iron-56 : 91.72% 55.93Iron-57 : 2.2% 56.93Iron-58 : 0.28% 57.93Before they start bead activity (next slide) give for HW The Atom and Its Properties WS 2 – Atomic MassAnd give out The Atom and Its Properties WS 3 – Ch. 5.1 and 5.2 Definitions
• Make sure you note which bag has which abundance.Set up a table w/#p+ and #n0 as well as mass number. For simplicity, we’ll consider each proton and each neutron to be 1 amu. Submit as group activity.
• Go over how to read this chart. Wavelength in meters is on top. Frequency in Hertz is on the bottom.Take-home message of this slideThe higher the frequency (or the shorter the wavelength) the more energy the light has. So, UV light has more energy than IR light. (It is the UV light causes sunburn.)The ‘Do Now’ I have for this has problems like “Determine the type of radiation with a wavelength of 1 m.” - Answer: RadioGo over a couple of more of these type of questions (and frequency). Good for “Do Now” and test.
• I’m pushing this a little more than last year. Keep equations and math simple. Save complicated ones for bonus questions.
• Use an example of throwing something into a swimming pool: With large items K.E. is large and water (like the electrons) will “splash out”. With smaller energy items like a “pin” water (like the electrons) will not splash out.Light has to be a certain frequency to eject the electron from the metal. The intensity (amplitude) doesn’t matter.
• Start of Chapter 5.2Ch. 5.2VocabGround StateDe Broglie EquationHeisenberg Uncertainty PrincipleQuantum MechanicsMight use the POGIL packet in folder in conjunction with this along with the salts and H2 gas demo.
• Hydrogen Spectrum.First seen in the mid-1800’s.
• The n is the same as Balmer’s n. Only this time, Bohr defined it as something rather than just an integer.
• The n is the same as Balmer’s n. Only this time, Bohr defined it as something rather than just an integer.
• Principal quantum number n=1 is closest to nucleus; n=7 is farther awayWhen all electrons are in their lowest state it is called the “ground state”When electrons are hit with energy (thermal, electrical etc.) they may go to an “excited state” atoms have multiple excited statesBohr model only worked for hydrogen, but it did a great job of predicting electronic transitions.Each circle represents an allowable energy state.When the electron goes back to ground state, a photon is emitted.
• Hydrogen Spectrum.Show actual hydrogen and other gases using tubes and power supply.
• Call the sublevels “l” (the letter ell) this year. This is so I can be more comfortable in assigning chapters 6 &amp; 7 for the AP summer assignment.After this slide, give out The Atom and Its Properties WS 5 – Electron OrbitalsAnd The Atom and Its Properties WS 6 – Ch. 5.3 Definitions
• Start of Ch. 5.3Ch. 5.3 VocabElectron configurationAufbau principlePauli Exclusion PrincipleHund’s RuleValence ElectronElectron Dot Structure
• Link is to Electron Configurations video
• This shows a summary of what we just did with the interactive periodic table.
• 6.1 VocabularyPeriodic LawGroupPeriodRepresentative ElementTransition ElementMetalAlkali MetalAlkali Earth MetalTransition MetalInner Transition MetalNonmetalHalogenNoble GasMetalloid
• Vocab given in The Atom and Its Properties W S 8 – Ch. 6 Vocab
• Lavoisier was on the wrong side of history. He associated with the aristocracy in France during the French Revolution and got guillotined.
• Italian chemist StanislaoCannizzaro published a list of atomic weights used at the First International Chemical Conference in 1860.
• Remember that protons, neutrons and electrons were not discovered until the late 19th and early 20th centuries.Mendeleev used Cannizzaro’s list to start his work.
• The numbers above represent atomic masses.
• Note: Number of valence electrons for Representative Group is either Group # (Using A &amp; B notation) orGroup # (for Groups 1 &amp; 2) or (Group # - 10) for the 1 – 18 Group # notation.Electron dot notation is just used for the Representative elements since the valence electrons are constant. The transition elements can vary how many electrons are actually used.
• These are Problems #7 and 8 on p. 162. (I added the last one in the chart to have a D-block example.)Config Group Period Block[Ne]3s2 2 (2A) 3 S[He]2s2 2 (2A) 2 S[Kr]5s24d105p5 17(7A) 5 P[Ar]4s23d5 7 (7B) 4 DAfter this slide give out The Atom and Its Properties WS 9 – Periodic Table Basics as CW/HWAnd The Atom and Its Properties WS 10 – Ch. 6.3 Definitions
• 1.Mg: [Ne] 3s2  Mg2+: [Ne] 2. O: [He]2s22p4  O2-: [He]2s22p6 or [Ne]
• P. 175 – Problem #63 for practice.Which element in each pair has the larger 1st ionization energy?Na or Al  Al why? Smaller atom in same period.Ar or Xe Ar why? Smaller atom in the same group.Ba or Mg  Mg why? Smaller atom the same group.
• Can show test tubes of carbon, silicon, tin, lead.Can even show ‘oxygen’, sulfur, maybe selenium?
• This is Exercise #22 on p. 169 of text.Answers: F, Br, Br, F, Br (it’s a liquid instead of a gas, molecules closer together)After this slide, give out The Atom and Its Properties WS 11 – Periodic Trends
• ### Chapters 4,5,6

1. 1. The Atom and ItsProperties Chapter 4 – Nucleus Chapter 5 – Electron Configuration Chapter 6 - Periodic Table
2. 2. Chapter 4 Objectives• Describe an atom‟s structure and differentiate among the particles that make it up.• Identify the numbers associated with elements and explain their meaning .• Realize that the number of protons in a nucleus defines an element.• Calculate the average atomic mass given isotopes and relative abundance2
3. 3. Chapter 4 VocabularyChapter 4.1 Chapter 4.3 • Dalton‟s Atomic • Atomic Number Theory • Isotope • Atom • Mass NumberChapter 4.2 • AMU (Atomic Mass Unit) • Electron • Average Atomic Mass • Nucleus • Proton • Neutron 3
4. 4. Brief History • Many ancient scholars believed matter was composed of such things as earth, water, air, and fire. • Many believed matter could be endlessly divided into smaller and smaller pieces.
5. 5. Brief History • Democritus (ancient Greek philosopher in about 460-370 BCE) believed matter was made up of tiny particles, „atomos‟ in Greek, which gave us our modern word atom. • In other words, matter could not be infinitely divided.5
6. 6. Brief History• Democritus had many other ideas that were left unexplored until school teacher John Dalton (1766 – 1844) revived them in 1803 to coincide with the industrial revolution.6
7. 7. Dalton’s Atomic Theory - 18031. Each element is composed of atoms.2. Atoms of a given element are identical. (Atoms of a specific element are different from atoms of another element.)3. Atoms cannot be created or destroyed, only rearranged, combined or separated.4. Different atoms combine in simple whole number ratios to form compounds.7
8. 8. Definitions• Electrons (e-)are negatively charged particles. They were the first subatomic particle discovered.• Protons (p+) are positively charged particles found in the nucleus. Discovered by Ernest Rutherford (1871 – 1937) in 1911. (More on him in the movie)• Neutrons (n ) are subatomic particles about the size of a proton but carry no charge. Discovered by James Chadwick (1891-1974) in 1932.8
9. 9. Modern View of the AtomThe nucleus is where the protons and neutronsare located and contain most of the atom‟smass.9
10. 10. Protons, Neutrons and Electrons Particle Symbol Charge Relative Mass Electron e- 1- 1/1840 Proton p+ 1+ 1 Neutron n0 0 ~110
11. 11. How Atoms Differ – Ch. 4.3 Mass # = #p+ + #n C-12 C-1411
12. 12. •Sometimes Atomic Symbols areDisplayed as:12
13. 13. •Displayed on Periodic Table13
14. 14. Mass Number• Mass Number is the number of protons plus the number of neutrons of a particular isotope of an element. (Mass # = #p+ + #n )• Thus we have atoms called: • Potassium -39 • Potassium -40 • Potassium -41• What is the number of neutrons in each of these isotopes?14
15. 15. Isotope ExamplesIsotope Atomic Mass Number Number Number Number Number Protons Neutron Electron 52Cr 24 52 24 28 24 75As 33 75 33 42 33222Rn 86 222 86 136 8615
16. 16. What’s all this amu business?• To simplify a system of indicating atomic masses since protons and neutrons have such extremely small masses, scientists have assigned the carbon-12 atom a mass of exactly 12 atomic mass units. (amu)• The mass of 1 amu (1/12 the mass of carbon- 12) is very nearly equal to the mass of a single proton or neutron but not the same.• 1 amu = 1.66 x 10-24 grams16
17. 17. Weighed Averages• Your QPA is a weighted average. The more credit a course is worth, the more an “A” will help your grade, and a “D” will hurt it.• Let‟s look at the example in the handbook.17
18. 18. •Your total number of credits is 41 and total quality pointsis 3667.This makes your QPA average 3667/41 = 89.44 18
19. 19. Atomic Mass• Why is potassium‟s atomic mass 39.098 in the periodic table?• How about • Atomic mass for Cl is 35.453? • Or atomic mass for Li at 6.941?• The atomic mass is the weighted average of mass numbers of the isotopes.• Based on abundance of each isotope.19
20. 20. •Isotopes and Mass Number• Your text (p. 119) shows how to calculate the mass number for Cl given the % abundance of the isotopes.• Let‟s do this for another element: Li • 6Li is 7.59 % abundant; 6.015 amu • 7Li is 92.41% abundant; 7.015 amu• Method 1: Use percentages. Think of this as a sample of 100 atoms.20
21. 21. •In Tabular Form Species Mass Abundance Mass x (amu) % Abundance (Weighted share) 6Li 6.015 7.59 45.65 (isotope) 7Li 7.015 92.41 648.26 (isotope) Li 100.00 694.41 (100 atoms) Li (atom) 6.9421
22. 22. Average Atomic Mass• The average mass of an atom is found by weighting the natural abundances of its isotopes.• Lithium (Method 2): Change % to fraction. • 6Li 6.015 amu 7.59% = 0.0759 • 7Li 7.015 amu 92.41% = 0.9241 Mass (amu) Frac abund Mass share Avg mass = 6.015 amu x 0.0759 = 0.46 amu 7.015 amu x 0.9241 = 6.48 amu 6.94 amu/atom22
23. 23. 23
24. 24. Bead Activity• Note what set you have.• Count the number of protons (red beads) and neutrons (green beads) each isotope (bag) has.• Use the abundance listed on each bag to calculate the weighted share.• Use the weighted share of each isotope to determine the average atomic mass.• Also, determine which element you have.24
25. 25. •Set #_______ Names:________________ Mass x Mass # Abun- Abundance #p+ #n0 (Isotope dance mass) % (Weighted share) Bag 1 Bag 2 What is the element? What is its average atomic mass?25
26. 26. Electrons in AtomsChapter 5
27. 27. •Chapter 5 Objectives • Compare wave and particle matters of light • See how frequency of light emitted by an atom is unique to that atom • Compare and contrast the Bohr and quantum mechanical models of the atom • Express the arrangements of electrons in atoms through orbital notations, electron configurations, and electron dot structures27
28. 28. Chapter 5 VocabularyChapter 5.1 Chapter 5.2• Electromagnetic Radiation • Ground State• Wavelength • Quantum Number• Frequency • Quantum Mechanical• Amplitude Model of the Atom • Atomic Orbital• Electromagnetic Spectrum • Principal Quantum• Quantum Number• Photoelectric Effect • Principal Energy• Photon Level• Atomic Emission Spectrum • Energy Sublevel28
29. 29. •Wave Nature of Light• Electromagnetic radiation displays wavelike behavior as it travels through space• Waves can be described by several common characteristics29
30. 30. •Characteristics of a Wave• Waves transfer energy• Properties of waves: • Frequency (ν – Frequency is the pronounced ‘nu’) - number of waves that hit this point in one Amplitude Number of vibrations second. per unit time – Hz λ (cycle/second) • Wavelength (λ) - Distance between 1.5 points on two 1 consecutive waves 0.5 0 • Speed of wave is -0.5 0 200 400 600 800 1000 1200 Frequency x wavelength -1 -1.5 Speed = ν x λ 30
31. 31. •Electromagnetic Spectrum Note: All EM Radiation travels at 3.00 x 108 m/s31
32. 32. •Electromagnetic Spectrum The speed of light (3.00 108 m/s) is the product of it‟s wavelength and frequency c = λν.
33. 33. The Electromagnetic Spectrum– all light is energy 33
34. 34. •Electromagnetic Spectrum• Gamma Rays – Highest frequency, shortest wavelength. Can pass through most substances• X Rays – Lower frequency than Gamma rays. Can pass through soft body tissue but can‟t pass through bone.• Ultraviolet (UV) Rays – Part of sunlight that causes sunburn
35. 35. •Electromagnetic Spectrum• VisibleLight – Sensitive to our eyes. Allows us to see color• Infrared – Less energy and longer wavelength than visible light. Felt as heat given off a heater or near a fire• Radio Waves – Lowest frequencies on the EM spectrum. Used by radio and over-the-air TV.
36. 36. An electromagnetic wave has a frequency of 6.0 x 104 Hz.Convert this frequency into its corresponding wavelength.Does this frequency fall in the visible region?x =c = c/ = 3.00 x 108 m/s 6.0 x 104 /s = 5.0 x 103 m No, it‟s a radio wave (~103 meters) 36 7.1
37. 37. •Examples Problems Answers • What is the C = λ*ν = Hint: nm frequency of green nanometer which is ν = C/ λ light, which has a 10-9 meters 8 m/s ν = 3.00 x 10 wavelength of 520 x 10-9 m 520 nm. = 5.77 x 1014 /s • A radio station λ = = mega which is M C/ν broadcasts at 8 94.7 MHz. What is 1063.00 x 10 m/s = Hz the wavelength of 94.7 x 106 /s the broadcast? = 3.17 m37
38. 38. •Nature of Light • Max Planck (1858- 1947) studied the different light emitted from heated objects • Matter can only gain or lose energy in small specific amounts38
39. 39. •Nature of Light • A quantum is the minimum amount of energy that can be gained or lost by an atom • The energy of EM radiation is proportional to its frequency (E α ν)39
40. 40. •Photons • Albert Einstein (1879- 1955) proposed that while a beam of light had wavelike characteristics, it also can be thought of as a stream of tiny particles (or bundles of energy) called photons • Each photon carries a quantum of energy40
41. 41. •Particle Nature of Light The photoelectric effect is when electrons are emitted from a metal‟s surface when light of a certain frequency shines on it.
42. 42. Ch. 5.2 – Quantum Theory of the Atom Each element has only certain specific frequencies of light that are emitted when atoms absorb energy and become excited Where do we see this? fireworks neon signs stars43
43. 43. •Hydrogen Spectrum44
44. 44. •Balmer Plot• In 1885, Johann Balmer observed the lines of the spectrum fit this surprisingly simple formula: 1 1 1 RH 2 2 n1 n2• Where n1 =2 and n2 = 3, 4, 5, etc.45
45. 45. •Balmer Plot 2500000 2400000 2300000 y = 1.0972E+07x + 4.0238E+02 2200000 R² = 1.0000E+00 1/Labmda, m-1 2100000 2000000 1900000 1800000 1700000 1600000 1500000 0.12 0.14 0.16 0.18 0.2 0.22 0.24 1/2^2 - 1/n^2 RH is the slope of this line, 1.0972 x 107 m-146
46. 46. •Electronic Energy Transitions • Neils Bohr (1885- 1962) proposed the model the hydrogen atom (1913) to explain the discreet nature of the hydrogen spectrum.47
47. 47. •Electronic Energy Transitions • Neils Bohr‟s model the atom (1913) • Electrons exist only in discrete, “allowable” energy levels • Energy is involved in moving electrons from one energy level to another • Principal quantum number (n) - specifies the electron‟s major energy level • The lowest energy is n=1, the next lowest in n=2, etc.48
48. 48. •Bohr’s Model of the Atom (cont’d) Bohr suggested that an electron moves around the nucleus in only certain allowed circular orbits. n=2 n=150
49. 49. •Energy Absorption/Emission51
50. 50. 52
51. 51. •Atomic Emission Spectra53
52. 52. •Origin of Line Spectra Balmer series54
53. 53. 55
54. 54. •Quantum or Wave Mechanics Schrodinger applied idea of e- behaving as a wave to the problem of electrons in atoms. E. He developed the Schrodinger 1887-1961 WAVE EQUATION Solution gives set of math expressions56 called WAVE
55. 55. •Wave Function• Thewave function predicts a three- dimensional region around the nucleus called the atomic orbital.58
56. 56. Waves •Wave motion: wave length and nodes “Quantization” in a standing wave59
57. 57. •Hydrogen Atom Solution Where: a0 is the Bohr Radius given by a0 = 4πεoh2/me2 Generalized Laguerre Polynomial m here is quantum number Constant = 2.18 x 10-18 J m is mass of electron60
58. 58. 61
59. 59. •Atomic Orbitals-Hydrogen62
60. 60. •Orbitals • No more than 2 e- assigned to an orbital • Orbitals grouped in s, p, d (and f) sublevels s orbitals p orbitals d orbitals63
61. 61. s orbitals p orbitals d orbitals s orbitals p orbitals d orbitals No. orbs. 1 3 5 No. e- 2 6 1064
62. 62. Energy Levels and Sublevels • Sublevels are grouped in energy level. • Each energy level has a number called the PRINCIPAL QUANTUM NUMBER, n which indicates relative size and energy of the orbitals • Row on PT indicates n65
63. 63. •Energy Levels and Sublevels n=1 n=2 n=3 n=466
64. 64. •QUANTUM NUMBERS The shape, size, and energy of each orbital is a function of quantum numbers: n (principal)  Energy Level l (angular)  sublevel (s, p, d, or f) which is its shape • Note: There are other quantum numbers that we will NOT discuss in detail. The „n‟ and „l‟ are sufficient.67
65. 65. QUANTUM NUMBERS Symbol Values Description n (principal) 1, 2, 3, .. Orbital size and energy level l (angular) 0, 1, 2, 3, … n-1 Orbital shape More commonly noted as: s, p, d, f, …n-1 or type (energy sublevel)68
66. 66. •Types of Atomic Orbitals70
67. 67. Types of Atomic Orbitals71
68. 68. •s Orbitals— Always Spherical Surface Surface of Dot density 90% picture 4πr2 versus probability distance sphere of electron cloud in 1sSee Active Figure 6.1372
69. 69. •p Orbitals The three p orbitals lie 90o apart in space73
70. 70. •2px Orbital 3px Orbital PLAY MOVIE74
71. 71. •Hydrogen-like Orbitals (at most two electrons/orbital) n Sublevel Orbitals Max. Max (l) Orbital Elec n2 2n2 1 s s 1 2 2 s s px, py, pz 4 8 p 3 s s p px, py, pz 9 18 dxy,dxz,dyz,dx2-y2, dz2 d 4 s s p px, py, pz dxy,dxz,dyz,dx2-y2, dz2 16 32 d And 7 f orbitals75 f
72. 72. Chapter 5.3 Vocabulary • Electron configuration • Aufbau principle • Pauli Exclusion Principle • Hund‟s Rule • Valence Electron • Electron Dot Structure76
73. 73. •Electron Configurations • An atom‟s electron configuration is the arrangements of electrons in the atom. • Electrons are arranged to minimize energy. • In other words, electrons fill up the lowest energies possible first. This is the Aufbau Principle.77
74. 74. •Assigning Electrons to Atoms • Electrons generally assigned to orbitals of successively higher energy. • For H atoms, E depends only on n. • For many-electron atoms, energy depends on both n and l.78
75. 75. •Energy Level Diagram of Hydrogen79
76. 76. •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 n + l increases. b) for subshells of same n + l, subshell with lower n is PLAY MOVIE lower in energy.80
77. 77. Orbitals and Their EnergiesMany-Electron Atoms 81
78. 78. •Aufbau Diagram -- Filling ElectronOrbitals 1s 2s 3s 4s 5s 6s 7s 8s Start here 2p 3p 4p 5p 6p 7p n+l=1 n+l=2 n+l=3 n+l=4 3d 4d 5d 6d Haven‟t gotten this far. What orbitals are being filled with n+l=5 elements 110-118? n+l=6 4f 5f n+l=7 n+l=8The orbital with the lower „n‟ is lower in energy if the n+l number is the same.82
79. 79. •Writing Atomic Electron Configurations Two ways of writing spdf notation configs. for H, atomic number One is 1 no. of called 1s electrons the spdf notation value of l value of n .83
80. 80. •Pauli Exclusion Principle No two electrons in the same orbital can have the same spin. One electron is spin up, the other is spin down.84
81. 81. •Writing Atomic Electron Configurations Two ways of writing configs. ORBITAL BOX NOTATION for He, atomic number = 2 Other is called Arrows the orbital box 2 1s depict notation. electron 1s spin It would be a violation of the Pauli exclusion principle to have both of these electrons as spin up or both as spin down.85
82. 82. •Electron Configurations and the Periodic Table86
83. 83. •Lithium Group 1 (1A) Atomic number = 3 1s22s1  3 3p total electrons 3s Interactive Periodic Table 2p 2s 1s87
84. 84. Ground State Electron Configurations88
85. 85. •Carbon Group 14 (4A) Atomic number = 6 1s2 2s2 2p2  3p 6 see for the first time total 3s Here we electrons HUND’S RULE. When 2p placing electrons in a set of 2s orbitals having the same energy, we place them singly as 1s long as possible.89
86. 86. Electron Configuration for Elements 11-18 Noble gas notation uses noble gas symbols in brackets to shorten inner electron configurations of other elements.90
87. 87. Sodium Group 1 (1A) Atomic number = 11 1s2 2s2 2p6 3s1 or “neon core” + 3s1 [Ne] 3s1 (uses noble gas notation) Note that we have begun a new period. All Group 1A elements have [core]ns1 configurations.91
88. 88. Electron Configurations and the Periodic Table92
89. 89. Transition Metals All 4th period elements have the configuration [argon] nsx (n - 1)dy and so are d-block elements. Chromium Iron Copper93
90. 90. Transition Element Configurations 3d orbitals used for Sc-Zn94
91. 91. 95
92. 92. Electrons in Energy Levels• Electrons fill up levels from lowest energy to highest energy (Aufbau Principle)• Outermost electrons are called Valence Electrons.• When atoms come close together it is the Valence Electrons that interact.96
93. 93. Valence Electrons • Howto determine which electrons are in outer shell? • Write down electron configuration of an element in noble-gas configuration • Whatever electrons are displayed in the highest energy shell (n) only are valence electrons (Main Group elements)97
94. 94. Lewis Dot Diagrams • How do we represent Valence Electrons? • By a Lewis Dot Diagram • Rules for Lewis Dot Diagrams: • Use the Elemental Symbol • Use 1 dot to represent each valence electron • The symbol represents the nucleus and all the inner (core) electrons. • Examples: . .. .. • Li·, Be: , ·C·, ·Cl:, :Ne: . .. ..98
95. 95. Valence Electrons • Examples • O given by [He]2s22p4 so it has 6 valence electrons. • O •• •• • • Ga given by [Ar]3d104s24p1 has 3 valence electrons. •Ga• •99
96. 96. •Electron Dot Representation100
97. 97. Chapter 6 101
98. 98.  Draw the periodic table and label the electron blocks and areas of non-metals, metals, and metalloids. Relate the Lewis dot structure to its place in the periodic table. Explain periodic trends as one moves along periods and down groups in the periodic table 102
99. 99. Chapter 6.1-6.2  TransitionMetal Periodic Law  Inner Transition Group Metal Period  Lanthanide Series Representative  Actinide Series Element  Nonmetal Transition Element  Halogen Metal  Noble Gas Alkali Metal  Metalloid Alkaline Earth Metal 103
100. 100.  Antoine Lavoisier (1743-1794) in the 1790s compiled a list of 23 known elements.  Many known since ancient times (copper, gold, silver, etc)  Others had been more recently discovered (oxygen, hydrogen, carbon, an d others) 104
101. 101. Scientists in 1860agreed upon a commonmethod to determineelemental mass firstpublished by StanislaoCannizzaro (1826-1910). 105
102. 102.  Earlyattempts organized the elements by increasing mass  The properties were roughly periodic, but some elements were out of order  Mendeleev’s Table (ca. 1870) was the most notable effort at organizing the elemental properties 106
103. 103. Dmitri Mendeleevnoticed in his tablethat there wererepetitions ofphysical andchemical propertieswhen the elementswere arranged byatomic mass. 107
104. 104. 108
105. 105. Properties of Germanium (Ge)Property Predicted (1869) Actual (1886)Atomic Mass 72 u 72.6 uColor Dark gray Gray-whiteDensity 5.5 g/mL 5.32 g/mLMelting Point High 937⁰ CDensity of Oxide 4.7 g/mL 4.70 g/mLOxide solubility in Slightly dissolved Not dissolvedHCl by HCl by HClFormula of chloride EsCl4 GeCl4 109
106. 106.  PeriodicLaw states that chemical and physical properties repeat in regular cyclic patterns when they are arranged by increasing atomic number.  Startswith metals at left and goes to non-metal (noble gas) on right  Properties change in orderly progression across a period. 110
107. 107. 111
108. 108. Columns, Noble Periodic Table Alkali Metals Groups or Halogens Gases Families Alkaline Earth Metals Representative Elements Transition ElementsPeriods Inner Transition Elements Metals Metalloids Nonmetals 112
109. 109.  What are some of the elemental properties that make the periodic table, well, periodic? Classification by metals, nonmetals and metalloids  Metals - shiny ductile, malleable solids, good conductors of heat and electricity  Nonmetals - dull, brittle solids; or gas, poor conductors of heat and electricity  Metalloids - have chemical and physical properties of both metals and nonmetals 113
110. 110.  Representative Elements (Sometimes called A Group)  Group # = number of valence electrons  Means similar Lewis dot structure and similar properties.  s-block elements have 1-2 electrons in s-orbital  p-block elements have 1-6 electrons in p-orbitals  Noble gases have filled valence shells  Energylevel of valence electrons is at energy level given by period (row) number 114
111. 111.  Transition Elements (Sometimes called B Group)  d-block elements have 1-10 electrons in d- orbitals  Columns 3-12 in periodic table  Energy level of valence electrons at n and partially filled n-1 d orbitals (example: 4s and 3d)  f-block (Lanthanides and Actinides) have 1-14 electrons in f-orbitals 115
112. 112. 1 2 13 14 15 16 17 18 116
113. 113.  Fill in the missing info for the following elements: Configuration Group Period Block [Ne]3s2 2 (2A) 3 S [He]2s2 2 (2A) 2 S [Kr]5s24d105p5 17 (7A) 5 P [Ar]4s23d5 7 (7B) 4 D Identify the element fitting the description. a) Group 2 (2A) element in 4th period: Calcium b) Noble gas in 5th period: Xenon c) Group 12 (2B) element in 4th period: Zinc d) Group 16 (6A) element in 2nd period: Oxygen 117
114. 114. PLAY MOVIE PLAY MOVIE PLAY MOVIE 118
115. 115.  Effective Nuclear Charge (Z*) – Not in book! Shielding Ion IonizationEnergy Octet Rule Electronegativity 119
116. 116.  Atomic and ionic size  Ionization energy  Electronegativity  Metallic Character Higher effective nuclear charge Electrons held more tightlyLarger orbitals.Electrons held lesstightly. 120
117. 117.  Z* is the nuclear charge experienced by the outermost electrons. (Note: not in book!) Z* increases across a period owing to shielding by inner electrons. Shielding is blocking by inner electrons. For a period (row), the number of shielding electrons remain the same, but the number of protons in the nucleus increases. Example: All elements in the 121
118. 118.  So we can estimate as Z* = [ Z - (no. inner electrons) ] or Z* = Z – S (inner electrons) Z is total number of electrons S is the number of electrons blocking the valence shell electrons, the underlying noble gas electrons. Charge felt by 2s e- in Li 122
119. 119. Orbital energies “drop” as Z* increases 123
120. 120.  Atomicsize is a periodic trend influenced by electron configuration. For metals, atomic radius is half the distance between adjacent nuclei in a crystal of the element. 124
121. 121.  Forother elements, the atomic radius is half the distance between nuclei of identical atoms that are bonded together. 125
122. 122. 126
123. 123.  Size (radius) goes UP on going down a group. See previous slide. Because electrons are further from the nucleus, there is less attraction. Size (radius) goes DOWN on going across a period. 127
124. 124. Size (radius) decreases across a period owing to increase in Z*. Each added electron feels a greater and greater positive charge.Note: Electrons in the same energy level don’t shield each other too much. Large Small Increase in Z* 128
125. 125. Radius (pm)250 K 3rd period 1st transition200 series 2nd period Na Li150 Kr100 Ar Ne 50 He 0 0 5 10 15 20 25 30 35 40 Atomic Number 129
126. 126.  The radius of an atom when it has become an ion. An ion is an atom or bonded group of atoms that has a positive or negative charge. An atom acquires a positive charge by losing electrons or negative charge by gaining electrons!! 130
127. 127. To form positive ions from elements remove 1or more e- from subshell of highest n [orhighest (n + l)].Al: [Ne] 3s2 3p1 - 3e-  Al3+: [Ne] 3s0 3p0 3p 3p 3s 3s 2p 2p 2s 2s 1s 1s 131
128. 128. Atoms tend to gain, lose, or share electrons to get 8 valence electrons(except small atoms up to Boron) 132
129. 129. 1. Write the electron configuration and orbital box diagram for Mg when it is an ion. Hints: What is its noble gas configuration? What will they do to get an octet?2. Write the electron configuration and orbital box diagram for O when it is an ion. 133
130. 130. + Forming aLi,152 pm + Li , 78 pm positive3e and 3p 2e and 3 p ion. Positiveions are SMALLER than the atoms from which they come. The electron/proton attraction has gone UP and so size DECREASES. Electron Configuration as ion is: [He] 2s0 134
131. 131. - Forming aF, 71 pm F- 133 pm negative ,9e and 9p 10 e and 9 ion. p  Negative ions 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.  Electron configuration as ion: 1s22s22p6 (just like neon.) 135
132. 132. See Figure 6-14 136
133. 133. Why do metals lose electrons in their reactions?Why does Mg form Mg2+ ions and not Mg3+?Why do nonmetals take on electrons? 137
134. 134. IE = energy required to remove an electron from an atom in the gas phase. PLAY MOVIE Mg (g) + 738 kJ  Mg+ (g) + e- 138
135. 135. IE = energy required to remove an electron from an atom in the gas phase. Mg (g) + 738 kJ  Mg+ (g) + e- PLAY MOVIE Mg+ (g) + 1451 kJ  Mg2+ (g) + e-Mg+ has 12 protons and only 11 electrons. Therefore, IE for Mg+ > Mg. 139
136. 136. 1st: Mg (g) + 735 kJ  Mg+ (g) + e- 2nd: Mg+ (g) + 1451 kJ  Mg2+ (g) + e- PLAY MOVIE3rd: Mg2+ (g) + 7733 kJ  Mg3+ (g) + e-Energy cost is very high to dip into a shellof lower n. 140
137. 137. 141
138. 138. 1st Ionization energy (kJ/mol)2500 He Ne2000 Ar1500 Kr1000 500 0 1 3 5 7 9 11 13 15 17 19 21 23 25 27 29 31 33 35 H Li Na K Atomic Number 142
139. 139. As Z* increases, orbital energies “drop” and IE increases. 143
140. 140. 144
141. 141.  IE increases across a period because Z* increases.  Metals lose electrons more easily than nonmetals.  Nonmetals lose electrons with difficulty.High ionization energy: atoms wantto hold on to electrons; likely to formnegative ionLow ionization energy: atom gives upelectron easily; likely to form positiveion 145
142. 142.  IEdecreases down a group Because size increases. Ability to lose electrons generally increases down the periodic table. See reactions of Li, Na, K 146
143. 143.  Which element in each pair has the larger 1st ionization energy? A. Na or Al B. Ar or Xe C. Ba or Mg 147
144. 144. PLAY MOVIELithium PLAY MOVIE PLAY MOVIESodium Potassium 148
145. 145. *Note: ‘metallic character’ not in book.An element with metallic character is one that loses electrons easily.Metallic character:• is more prevalent in metals on left side of periodic table• is less for nonmetals on right side of periodic table that do not lose electrons easily 149
146. 146. 150
147. 147.  Relativeability of an element to attract electrons in a chemical bond.  Ionization energy reflects ability of atom to attract electrons in an isolated atom  Generally, the higher the ionization energy of an atom, the more electronegative the atom will be in a molecule There are many electro negativity scales – we’ll use the one by Linus Pauling (values dimensionless) Will be used to determine things like polarity of a chemical bond. 151
148. 148. 152
149. 149. Decreases down a group  Why? Due to greater atomic radiusIncreases across a period  Why?Increased positive charge in nucleus (Greater Z*)Same trend as for ionization energy. Surprised? 153
150. 150.  Moving Left  Right (periods)  Z* Increases  Atomic & ionic Radius Decrease  Ionization Energy Increases  Electronegativity Increases  Metallic Character Decreases Moving Top  Bottom (groups)  Z* is roughly constant, but val e- distance increases  Atomic & Ionic Radius Increase  Ionization Energy Decreases  Electronegativity Decreases  Metallic Character Increases 154
151. 151. a) Electronegativity  Fluorineb) Ionic Radius  Brominec) Atomic Radius  Bromined) Ionization Energy  Fluorinee) Metallic character  Bromine 155