2. Brief Timeline of Atomic Theory

3. Democritus 400BC Greek philosopher

4. Hard Particle (Cannonball)Theory Proposed that they world was made up of tiny, indivisible particles moving through a void of empty space “atom” comes from the Greek word “atomos”, meaning indivisible (cannot be divided)

5. John Dalton 1808 AD First modern atomic theory

6. DaltonsDaltons Atomic Theory 1.All matter is composed of tiny, indivisible particles called atoms 2.All atoms of an element are identical 3.Atoms of different elements are all different 4.Atoms combine in simple ratios to form compounds

7. J.J. Thomson 1897-1904 “Plum Pudding Model” Cathode Ray tube experiment demo

9. Cathode Ray Tube Thompson showed that cathode rays (electrons) were composed of negatively charged particles that separated from the gas atoms inside the tube Significant because: this meant that atoms are not hard, indivisible particles. Atoms are composed of smaller “subatomic” particles

10. Thomson’s Plum Pudding Model The atom was a hard sphere that was positively charged with negatively charged electrons that “dotted” the atom like raisins in plum pudding

11. The discovery of radioactivity Henri Becquerel – 1896 – Discovered that uranium ore released rays that could expose photographic film

12. The discovery of radioactivity Marie & Pierre Curie – Extracted 2 new elements from uranium (U)ore: radium (Ra) and polonium (Po) Marie Curie

13. Ernest Rutherford Magnetic Field Experiment Was able to separate radioactive rays into 2 types: alpha (  ) & beta (  ) Determined that  rays were composed of helium nuclei (He +2 charge) Gold Foil Experiment (1911) Lead to discovery of the nucleus, as a positively charged center of atom, containing the mass Most of the atom is negatively charged empty space, electrons are outside the nucleus

14. Magnetic Field Experiment

16. Gold Foil ExperimentExperiment

17. Gold Foil Experiment

18. Experiment

19. Gold Foil Experiment

20. Rutherford’s Atomic Model

21. Rutherford’s “Nuclear Model” Most of the atom is negatively charged empty space, surrounding a small, positively charged nucleus, containing most of the mass of the atom

22. Modern Theory of Atomic Structure Developed by Niels Bohr, based on the science of nuclear physics Bohr determined that an element's position on the periodic table was related to its electron configuration.

23. Electron configuration Electron configuration – shows how many electrons are in each energy level or “ring” Ex: Carbon 2-4

24. Bohr’s Planetary Atomic Model Niels Bohr (1922) Determined that electrons rotate around the nucleus in discrete paths or rings

25. Planetary Model of Atomic Structure

26. Wave-Mechanical Model Current (modern) theory of atomic structure Moseley used x-ray analysis to calculate an integer for each element: these integers are the atomic numbers

27. Wave-Mechanical Model There is a tiny, dense positively charged nucleus at the center of a huge negatively charged electron cloud

28. Wave-Mechanical Model

29. Orbital Region of probability of finding an electron

30. “The whole point:” The modern model of the atom is the result of many investigations that have been revised over a long period of time by many scientists Atomic theory song

31. Place the models of atomic structure in order from earliest to the modern theory:

32. Basic Atomic Structure The nucleus occupies less than 0.01% of the total volume of an atom but accounts for 99.97% of its mass. Thus most of an atom is EMPTY SPACE where the ELECTRONS are found, this is called an ELECTRON CLOUD. One atomic mass unit is 1/12 TH THE MASS OF A CARBON-12 ATOM. This is the standard by which the masses of all other elements are determined. It is abbreviated “u”.

33. Subatomic Particles

35. Atomic NumberMass Number Nuclear Charge # of Protons # of Neutrons# of electrons 27 13 Al 35 17 Cl 11H11H 207 82 Pb Use your Periodic Table to complete the following:

36. The only number that never changes for an element is ATOMIC NUMBER !!

37. NameSymbol Atomic Number Mass Number Charge # of Protons # of Neutrons# of electrons 19 9 F 0 Helium-4 0 11230 Nitrogen-14 0 014 320 16 64 29 Cu 0 25 0 35 0 8156 531310 05374 Atomic Structure 1

38. Name-massSymbol Atomic Number Mass Number Charge # of Protons # of Neutrons# of electrons Flourine-19 19 9 F9 19 09 10 9 Helium-4 4 2 He 2 40 22 2 Sodium-23 23 11 Na1123011 1211 Nitrogen-14 14 7 N7 14 07 77 Silicon-28 28 14 Si14 28 014 Silicon-32 32 16 Si 16320 16 copper-64 64 29 Cu29 64 029 3529 Manganese-60 60 25 Mn2560 025 3525 Barium-137 137 56 Ba56 137 056 8156 Iodine-131 131 53 I531310537853 Iodine-127 127 53 I53 127 05374 53 Atomic Structure 1

39. Phosphorus-32 0 14 6 C 0 Potassium-39 0 160 8 56 26 Fe 0 18400 02935 791970 24 12 Mg 0 **Shade the columns representing the nucleons light blue

40. ISOTOPE Forms of the same element having different mass due to different number of neutrons. Indicated by “element name-mass”

41. 15 8 O 16 8 O Name: _______________ Mass: ________________ Protons: ______________ Neutrons: _____________ Name: _______________ Mass: ________________ Protons: ______________ Neutrons: _____________

42. Practice: NameSymbolAtomic #Mass ## Protons# Neutrons# Electrons 235 U 238 U Carbon-12 Carbon-13

43. The mass number on the periodic table indicates the weighted average of all the naturally occurring isotopes of an element To calculate a weighted average: % X mass + % X mass + …..100

44. Neon is naturally found in nature having 90.51% mass of 20.00u, 0.24% mass of 21.00u and 9.22% mass of 22.00u. Calculate the weighted atomic mass of neon.

46. Calculate the weighted average atomic mass of the elements below. Show all work, round to the nearest hundredth. a.)99.63% 14 N & 0.37% 10 N b.)69.1% 63 Cu (actual mass of 63.93g) & 30.9% 65 Cu (actual mass of 64.93g) c.)78.9% 24 Mg, 10.00% 25 Mg & 11.01% 26 Mg

47. You can estimate which isotope is found in the highest abundance as the one with a mass closest to the mass listed on the periodic table Example: Chlorine-35 mass 34.969g Chlorine-37 mass 36.966g Look on the periodic table for the mass of chlorine ____________________________ The more abundant isotope has a mass closer to the mass given on the periodic table:_____________

48. Practice: Which isotope of silicon would be found in the highest percentage? 28 14 Si, mass 27.977 29 14 Si, mass 28.976 30 14 Si, mass 29.974 Why?

49. Atomic Structure 2 IsotopicNotation Number of protons Number of neutrons Number of electrons Mass number 1.Oxygen-16O-16 16 O 2.Oxygen-18 3.Ar-40 4. 18 5. 16 32 6. 34 S 7. 1920 8. 19 41 9. Iron- 10. 57 Fe 11. 2632 12.Ne-20 13. 10 22 14.Hydrogen- 1 15.H-2 16. 3H3H

50. 2.) Calculate the weighted average of the following naturally occurring isotopes. SHOW ALL WORK! a.) 95.50% 7 Li & 7.50% 6 Li d.) 99.63% 14 N & 0.37% 15 N b.)80.20% 11 B & 19.80% 10 B e.) 78.9% 24 Mg, 10.00% 25 Mg, & 11.01% 26 Mg c.)95.02% 32 S, 0.75% 33 S, & 4.21% 34 Sf.) 92.23% 28 Si, 4.67% 29 Si, & 3.10% 30 Si

51. Changes in number of subatomic particles Isotopes Change in number of neutrons Same atomic number, different mass Same number protons, different number neutrons Ions Change in number of electrons A ca t ion is positive ion, results from loss of electrons, reducing radius An a n ion is negative ion, results from gain of electrons, increasing radius

52. IONS A charged part of an atom, resulting from the loss or gain of electrons VALENCE electrons: outermost electrons, the last number in an electron configuration KERNEL electrons: all electrons except valance electrons

53. Electron configuration Electron configuration – shows how many electrons are in each energy level or “ring” Ex: Carbon 2-4

54. Electron configuration of sodium:

55. 2 diagrams of atomic structure: Bohr diagramsLewis electron dot diagrams Bohr realized that the rows on the periodic table corresponded to the number of shells of electrons Lewis realized that the groups/families on the periodic table correspond to the number of valence electrons This model shows the nucleus, indicating the number of protons and neutrons, surrounded by rings, representing each energy level This model shows the element symbol surrounded by dots, representing the valence electrons. You must place one dot at each (3, 6,9,12 o’clock) location before “doubling up” (exception: Helium) 18 9 F electron configuration 2-7 F electron configuration 2-7

56. 118 1 Bohr Atomic Structures tables to fill in the electron 4 configurations, as shown, then draw the Bohr Atomic Structure for each element 1-20. 12 1213141516172 79111214161920 345678910 2-12-22-32-4 2324272831323540 1112131415161718 2-8-12-8-22-8-32-8-4 3940 Rules: 1.) Show placement2.) The nucleus is3.) Indicate the number of ALL electronsrepresented by a centerof electrons in each circle showing theenergy level, by writing *use atomic # # of protons & thethe number on each ring. 1920 OR the entire# of neutrons 2-8-8-12-8-8-2 electron configuration** closest to nucleus is 1 st

58. 2 Main Types of Ions: a n ion A negative ion Ex: Cl -, O -2 ca t ion A positive ion Ex: Na +, Al +3

59. The octet rule Atoms will gain or lose electrons in order to have a full valence shell of 8 electrons. Exception: Helium can have a maximum of 2 valance electrons

60. When an atom gains 1 or more electrons It becomes a negative ion and it’s radius increases. A negative ion is an anion.

61. When an atom loses 1 or more electrons It becomes a positive ion and it’s radius decreases. A positive ion is a cation.

63. CATIONANION Definition positive ionnegative ion Results from Loss of electron(s)Gain of electron(s) Indicated by (+) charge(-) charge What happens to radius??? Gets smallerGets bigger Na Na + Naming “Element name-ion”Change ending of element to “-ide” Lewis Dot Structure [Na] +.. [:. F. :] -

64. How to predict if an element will form an anion or cation: The “electron clock”: 8/0 71 62 53 4 # valance electrons

65. Atomic Structure 3: Predicting Ions Element Electron configuration Lewis dot structure of atom Lose or gain electrons? How many electrons lost or gained? Ionic Charge ** Lewis dot structure of ion Radius increase or decrease? F2-7Fgain1Fincrease Mg2-8-2Mglose2+2Mgdecrease O Al N Fr C

66. 2-8-8-1 2-8-7 2-8-18-18-8-2 2-8-6 2-8-5 2-3 **In the “ ionic charge ” column only: shade the cation charges red and the anion charges blue Element Electron configuration Lewis dot structure of atom Lose or gain electrons? How many electrons lost or gained? Ionic Charge ** Lewis dot structure of ion Radius increase or decrease?

67. # of Protons # of Neutrons # of Electrons Nuclear Charge Bohr Diagram of Atom Lewis Dot of Atom Predict Ionic Charge Lewis Dot of Ion Name of Ion ex 35 17 Cl 171817+17 Cl Chloride 1 23 11 Na 2 9 4 Be 3 65 30 Zn 4 14 7 N 5 32 16 S 6 20 10 Ne 7 127 53 I 8 108 47 Ag 9 70 31 Ga 10 12 6 C Atomic Structure 4

68. Atomic Spectra

69. Radiant Energy Energy that travels through space as electromagnetic waves at the speed of light

70. Electromagnetic Spectrum Includes all types of radiant energy from gamma rays (hi E) to radiowaves (lo E) Visible light is only a small portion of the spectrum

71. 1 photon = 1 quantum Quanta: tiny packets of energy released or absorbed by objects *Einstein and Plank determined that energy is released or absorbed in a continuous flow of small packets or quantum/photons

72. Release or Absorption of Energy:

74. Higher energy levels (excited state) Electrons absorb energy when jumping to Electrons release energy when falling to Lower energy levels (ground state)

76. Bohr used the emission spectrum as proof of planetary model But his model only works for hydrogen because he didn’t account for electrons moving between energy levels

77. Spectral Lines Characteristic wavelengths ( ) of photons of energy released as electrons fall from hi to lo energy

78. Spectral lines demo: Salt of Element Color of Flame Strontium Chloride Barium Chloride Copper (II) Chloride Lithium Chloride Potassium Chloride Identity Unknown Element Unknown Mixture

79. Emission Spectrum:

80. Each element has it’s own characteristic spectrum:

81. Compare H & He: hydrogenhelium

82. Because electrons do move between energy levels, emitting “spectral lines”, we had to change our view of atomic structure:

83. Excited State Electron Configurations Occurs when elements absorb energy and jump to a higher energy level. ** it will not look like it is written on periodic table, be sure they add to the correct number! Ground state: 2-8-1 Excited state : 2-7-2

88. “Crib Sheet” #p + = atomic number*#n 0 = mass-atomic number #e - = #p + - charge (use the sign of the charge) Isotope: same #p +, different #n o OR same atomic number, different mass To calculate weighted average: (%/100 x atomic mass) + (%/100 X atomic mass) + ….. *Ion: same # p +, different #e - Charge= #p + - #e -

89. Atomic Structure Review p. 17 1.11 2.9 3.43 4.92 5.118 6.13 7.11 8.4 9. Br 10. C 11. Sn 12. Zn 13. Cl 14. 40 15. 16

90. Atomic Structure Review p. 17 16.) =(.925x7) + (.0750x6) =6.475 +.45 =6.925 =6.93g 17) =(.789x24)+(.10x25)+(.1101x26) = 18.936 + 2.5 + 2.8626 = 24.2986 = 24.30g

91. Atomic Structure Review p. 18 18.) 2-8-1 19.) Na 20.) 2-7-2 21.) 19 22.) 1 23.) Y 24.) Ar 25.) Not possible 27.) as electrons fall from excited state to ground state energy is released as radiant energy (spectral lines). 28.) you can ID the gas element using spectral line analysis. 29.) electrons are negatively charged particles. B has 5 e -, its e - config. is 2-3, with 2 e - in the 1 st energy level and 3 e - in the 2 nd (valence) level

92. Atomic Structure Review MC?s 1.) 213.) 1 2.) 414.) 4 3.) 115.) 1 4.) 116.) 3 5.) 317.) 4 6.) 118.) 3 7.) 419.) 2 8.) 220.) 2 9.) 421.) 3 10.) 3 11.) 4 12.) 3pg 19-20 1.) 413.) 2 2.) 314.) 4 3.) 215.) 3 4.) 316.) 3 5.) 217.) 2 6.) 318.) 1 7.) 119.) 1 8.) 220.) 4 9.) 321.) 4 10.) 122.) 2 11.) 323.) 3 12.) 1pg 21-22

93. Atomic Structure Review p. 23 1.) 19p, 20n, 18e2.) 9p,10n,10e 3.) 5p,6n,2e4.) 15p,16n,18e 5.) 16p,16n,18e6.) 14p,14n,10e 7.) 7p,7n,10e8.) 20p,20n,20e 9.) 37p,48n,36e10.) 53p,75n,54e 11.) 30p,35n,28e12.) 6p,6n,10e