2. Atomic Structure Atoms and their structure Mr. Bruder

3. Dalton’s Atomic Theory n John Dalton (1766-1844) had four theories 1. All elements are composed of submicroscopic indivisible particles called atoms 2. Atoms of the same element are identical. The atoms of anyone element are different from those of any other element 3. Atoms of different elements can physically mix together or can chemically combine w/ one another in simple whole-number ratios to form compounds 4. Chemical reactions occur when atoms are separated, joined, or rearranged. However, atoms of one element are never changed into atoms of another elements as a result of a chemical reaction

4. Atoms & Subatomic Particles n Atom- smallest particle of an element that retains the properties of that element

5. n Gay-Lussac- under the same conditions of temperature and pressure, compounds always react in whole number ratios by volume. n Avagadro- interpreted that to mean n at the same temperature and pressure, equal volumes of gas contain the same number of particles n (called Avagadro’s hypothesis) A Helpful Observation

6. Electron n J.J Thomson (1856-1940) – discovered the electron in 1897 n Electron is the negative charged subatomic particle n An electron carries exactly one unit of negative charge & its mass is 1/1840 the mass of a hydrogen atom

7. Cathode Ray n The Cathode Ray tubes pass electricity through a gas that is contained at a very low pressure

8. Thomson’s Experiment Voltage source +-

9. Thomson’s Experiment Voltage source +-

10. Thomson’s Experiment Voltage source +-

11. n Passing an electric current makes a beam appear to move from the negative to the positive end Thomson’s Experiment Voltage source +-

12. n Passing an electric current makes a beam appear to move from the negative to the positive end Thomson’s Experiment Voltage source +-

13. n Passing an electric current makes a beam appear to move from the negative to the positive end Thomson’s Experiment Voltage source +-

14. n Passing an electric current makes a beam appear to move from the negative to the positive end Thomson’s Experiment Voltage source +-

15. Thomson’s Experiment n By adding an electric field

16. Voltage source Thomson’s Experiment n By adding an electric field + -

17. Voltage source Thomson’s Experiment n By adding an electric field + -

18. Voltage source Thomson’s Experiment n By adding an electric field + -

19. Voltage source Thomson’s Experiment n By adding an electric field + -

20. Voltage source Thomson’s Experiment n By adding an electric field + -

21. Voltage source Thomson’s Experiment n By adding an electric field he found that the moving pieces were negative + -

22. Thomson’s Atomic Model n Thomson’s Atomic Model n Thomson though electrons were like plums embedded in a positively charged “pudding”, so his model was called the “plum pudding” model

23. Thomsom’s Model n Found the electron n Couldn’t find positive (for a while) n Said the atom was like plum pudding n A bunch of positive stuff, with the electrons able to be removed

24. Mass of Electron n In 1909 Robert Millikan determined the mass of an electron with his Oil Drop Experiment n He determined the mass to be 9.109 x 10 -31 kg n The oil drop apparatus

25. Millikan’s Experiment Atomizer Microscope - + Oil

26. Millikan’s Experiment Oil Atomizer Microscope - + Oil droplets

27. Millikan’s Experiment X-rays X-rays give some drops a charge by knocking off electrons

28. Millikan’s Experiment +

29. They put an electric charge on the plates + + --

30. Millikan’s Experiment Some drops would hover + + --

31. Millikan’s Experiment + ++ +++++ -- -----

32. Measure the drop and find volume from 4/3πr 3 Find mass from M = D x V + + --

33. Millikan’s Experiment From the mass of the drop and the charge on the plates, he calculated the charge on an electron + + --

34. Proton n In 1886 Goldstein discovered the Proton n Proton is a positively charged subatomic particle found in the nucleus of a atom

35. Radioactivity n Discovered by accident n Bequerel n Three types –alpha- helium nucleus (+2 charge, large mass) –beta- high speed electron –gamma- high energy light

36. Ernest Rutherford n Rutherford (1871-1937) proposed that all mass and all positive charges are in a small concentrated region at the center of the atom n He used the Gold-Foil Experiment to prove his theory n In 1911 he discovered the Nucleus n Nucleus- central core of an atom, composed of protons and neutrons n The nucleus is a positively charged region and it is surrounded by electrons which occupy most of the volume of the atom

37. Rutherford’s Experiment n Used uranium to produce alpha particles n Aimed alpha particles at gold foil by drilling hole in lead block n Since the mass is evenly distributed in gold atoms alpha particles should go straight through. n Used gold foil because it could be made atoms thin

38. Lead block Uranium Gold Foil Florescent Screen

39. What he expected

40. Because

41. Because, he thought the mass was evenly distributed in the atom

42. What he got

43. How he explained it + n Atom is mostly empty n Small dense, positive piece at center n Alpha particles are deflected by it if they get close enough

44. +

45. Rutherford’s Model n Discovered dense positive piece at the center of the atom n Nucleus n Electrons moved around n Mostly empty space

46. Neutron n James Chadwick (1891-1974) – discovered the neutron in 1932 n Neutron is a subatomic particle with no charge but their mass is nearly equal to that of a proton

47. 46 Bohr Model n Bohr changed the Rutherford model and explained how the electrons travel. n Bohr explained the following in his model: 1. Electrons travel in definite orbits around the nucleus 2. Electrons are arranged in concentric circular paths or orbitals around the nucleus 3. Electrons don’t fall into the nucleus because electrons in particular path have fixed energy and don’t lose energy 4. His model was patterned after the motion of the planets around the sun. It is often called the Planetary model.

48. Bohr’s Model Nucleus Electron Orbit Energy Levels

49. 48 Quantum Theory n Bohr explained how electrons were moving via Quantum Theory n Key Terms: 1. Energy Levels- Regions around the nucleus where the electron is likely moving 2. Quantum- Amount of energy required to move an electron from one energy level to the next 3. Quantum Leap- Abrupt Change

50. Bohr’s Model Increasing energy Nucleus First Second Third Fourth Fifth } n Further away from the nucleus means more energy. n There is no “in between” energy n Energy Levels

51. 50 Bohr’s Model cont. n Energy levels are not equally spaced. n Energy levels more closely spaced further from the nucleus n Higher energy level occupied by an electron, the more energetic that electron is. n Amount of energy gained or lost by an electron is not always the same amount.

52. The Quantum Mechanical Model n Energy is quantized. It comes in chunks. n A quanta is the amount of energy needed to move from one energy level to another. n Since the energy of an atom is never “in between” there must be a quantum leap in energy. n Schrodinger derived an equation that described the energy and position of the electrons in an atom

53. Schrödinger’s Equation n The wave function is a F(x, y, z) n Actually F(r,θ,φ) n Solutions to the equation are called orbitals. n These are not Bohr orbits. n Each solution is tied to a certain energy n These are the energy levels Animation

54. n The atom is found inside a blurry “electron cloud” n A area where there is a chance of finding an electron. n Draw a line at 90 % The Quantum Mechanical Model

55. Modern View n The atom is mostly empty space n Two regions n Nucleus- protons and neutrons n Electron cloud- region where you have a chance of finding an electron

56. Quark n Protons & Neutrons can still be broken down into a smaller particle called the Quark n The Quark is held together by Gluons

57. Density and the Atom n Since most of the particles went through, it was mostly empty. n Because the pieces turned so much, the positive pieces were heavy. n Small volume, big mass, big density n This small dense positive area is the nucleus

58. Atomic Particles ParticleChargeMass (kg)Location Electron9.109 x 10 -31 Electron cloud Proton+11.673 x 10 -27 Nucleus Neutron01.675 x 10 -27 Nucleus

59. Subatomic particles Electron Proton Neutron NameSymbolCharge Relative mass Actual mass (g) e-e- p+p+ n0n0 +1 0 1/1840 1 1 9.11 x 10 -28 1.67 x 10 -24

60. Symbols n Contain the symbol of the element, the mass number and the atomic number X Mass number Atomic number

61. Sub-atomic Particles n Z - atomic number = number of protons determines type of atom n A - mass number = number of protons + neutrons n Number of protons = number of electrons if neutral

62. Symbols X A Z Na 23 11

63. Atomic Structure Symbols n Proton = p + n Electron = e - n Neutron = n 0 n Atomic # - Subscript n Mass # - Superscript

64. Rules for Atomic Structure 1. Atomic # = # of Protons 2. # of Protons = # of Electrons 3. Mass # = # of Protons + # of Neutrons n # of Neutrons = Mass # - # of Protons n If you know the Mass # & Atomic # you know the composition of the element

65. Symbols n Find n Find the –number –number of protons of neutrons of electrons –Atomic –Atomic number –Mass –Mass Number Br 80 35

66. Symbols n if an element has an atomic number of 34 and a mass number of 78 what is the –number of protons –number of neutrons –number of electrons –Complete symbol

67. Symbols n if an element has 78 electrons and 117 neutrons what is the –Atomic number –Mass number –number of protons –Complete symbol

68. Example Element Atomic # Mass #ProtonsElectro ns Neutro ns K19 115 1617 4623 35

69. Isotopes n Isotope- atoms that have the same number of protons but different number of neutrons n Since isotopes have a different number of neutrons the isotope has a different mass number. n Isotopes are still chemically alike because they have the same number of protons and electrons

70. Examples of Isotopes

71. © 2009, Prentice-Hall, Inc. Isotopes n Isotopes are atoms of the same element with different masses. n Isotopes have different numbers of neutrons. 11 6 C 12 6 C 13 6 C 14 6 C

72. Naming Isotopes n Put the mass number after the name of the element n carbon- 12 n carbon -14 n uranium-235

73. Electrical Charges n Electrical charges are carried by particles of matter n Atoms have no net electrical charges n Given the number of negative charges combines with the number of positive charges = Electrically Neutral n All elements are Electrically Neutral

74. Atomic Mass vs. Atomic Weight n Atomic Mass is for a single element n Most elements are Isotopes n How do we find their mass? n We use Atomic Weight

75. Measuring Atomic Mass n Unit is the Atomic Mass Unit (amu) n One twelfth the mass of a carbon-12 atom n Each isotope has its own atomic mass. We need the average from the percent abundance n Each isotope of an element has fixed mass and a natural % abundance n You need both of these values to find the Atomic Weight

76. Calculating Atomic Weight n Cl-35 34.969amu and 75.77% abundance n Cl-37 36.966amu and 24.23% abundance n To solve for Cl-35 1. AMU x Abundance 2. 34.969 x.7577 3. = 26.496 n You solve for Cl-37

77. Atomic Weight Cont. n Cl-37 1. AMU x Abundance 2. 36.966 x.2423 3. = 8.957 n Now you combine your two answers n 26.496 + 8.957= n 35.453 n Look at Cl on the table. What is the Atomic Weight?

78. Example n Calculate the atomic weight of copper. Copper has two isotopes. One has 69.1% and has a mass of 62.93 amu. The other has a mass of 64.93 amu. What is the atomic weight???

79. Atomic Weight & Decimals n Atomic Weight- of an element is a weighted average mass of the atoms in a naturally occurring sample of an element n Atomic Weights use decimal points because it is an average of an element