2. Atoms and Their Structure

3. History of the Atom n n Original idea (400 B.C.) came from Democritus, a Greek philosopher n n Democritus expressed the belief that all matter is composed of very small, indivisible particles, which he named atomos.

4. Who’s Next? n n John Dalton (1766- 1844), an English school teacher and chemist, studied the results of experiments by other scientists.

5. Dalton’s Atomic Theory   John Dalton proposed his atomic theory of matter in 1803.   Although his theory has been modified slightly to accommodate new discoveries, Dalton’s theory was so insightful that it has remained essentially intact up to the present time.

6. Dalton’s Atomic Theory 1. All matter is made of tiny indivisible particles called atoms. 2. Atoms of the same element are identical; those of different atoms are different.

7. Dalton’s Atomic Theory, cont. 3. Atoms of different elements combine in whole number ratios to form compounds 4. Chemical reactions involve the rearrangement of atoms. No new atoms are created or destroyed.

8. Parts of the Atom n n Because of Dalton’s atomic theory, most scientists in the 1800s believed that the atom was like a tiny solid ball that could not be broken up into parts. n n In 1897, a British physicist, J.J. Thomson, discovered that this solid- ball model was not accurate.

9. Parts of the Atom   Thomson’s experiments used a cathode ray tube.   It is a vacuum tube - all the air has been pumped out.

10. Thomson’s Experiment Voltage source +- Vacuum tube Metal Disks

11. Thomson’s Experiment Voltage source +-  At each end of the tube is a metal piece called an electrode, which is connected through the glass to a metal terminal outside the tube.

12. Thomson’s Experiment Voltage source +-  When the electrodes are charged, rays travel in the tube from the negative electrode, which is the cathode, to the positive electrode, the anode.

13. Thomson’s Experiment Voltage source +-  Because these rays originate at the cathode, they are called cathode rays.

14. Thomson’s Experiment Voltage source +-

15. Thomson’s Experiment Voltage source +-

16. Thomson’s Experiment Voltage source +-

17. Thomson’s Experiment Voltage source +-

18. Thomson’s Experiment By adding an electric field, By adding an electric field, + -

19. Voltage source Thomson’s Experiment + -

20. Voltage source Thomson’s Experiment n Thomson found that the rays bent toward a positively charged plate and away from a negatively charged plate. + -

21. Voltage source Thomson’s Experiment n He knew that objects with like charges repel each other, and objects with unlike charges attract each other. + -

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

23. Voltage source Thomson’s Experiment n J.J. Thomson concluded that cathode rays are made up of invisible, negatively charged particles referred to as electrons. + -

24. Cathode Ray Tube

25. Thomson’s Model   From Thomson’s experiments, scientists had to conclude that atoms were not just neutral spheres, but somehow were composed of electrically charged particles.

26. Thomson’s Model n n Matter is not negatively charged, so atoms can’t be negatively charged either. n n If atoms contained extremely light, negatively charged particles, then they must also contain positively charged particles — probably with a much greater mass than electrons.

27. Thomson’s Model n J.J. Thomson said the atom was like plum pudding, a popular English dessert. Draw Diagram in your notes Plum Pudding Model

28. Millikan’s Oil Drop Experiment n n R.A. Millikan found the charge of an electron to be -1.60 x 10 -19 Coulombs in his famous oil drop experiment.

29. The Proton   In 1886, scientists discovered that a cathode-ray tube emitted rays not only from the cathode but also from the positively charged anode.   Years later, scientists determined that the rays were composed of positively charged subatomic particles called protons.

30. The Proton   At this point, it seemed that atoms were made up of equal numbers of electrons and protons.

31. Ernest Rutherford n In 1909, a team of scientists led by Ernest Rutherford in England carried out the first of several important experiments that revealed an arrangement far different from the plum pudding model of the atom and discovered the nucleus.

33. Rutherford’s Experiment n n The experimenters set up a lead- shielded box containing radioactive polonium, which emitted a beam of positively charged subatomic particles through a small hole.

34. Rutherford’s Experiment n n The sheet of gold foil was surrounded by a screen coated with zinc sulfide, which glows when struck by the positively charged particles of the beam.

35. Lead block Polonium Gold Foil Florescent Screen

36. What Rutherford Expected n The alpha particles to pass through without changing direction very much.

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

38. the alpha particles should go straight through.

39. What Rutherford Observed

40. How Rutherford Explained It n n To explain the results of the experiment, Rutherford’s team proposed a new model of the atom. n n Because most of the particles passed through the foil, they concluded that the atom is mostly empty space.

41. How Rutherford Explained It  Alpha particles are deflected by the nucleus if they get close enough it.

42. How Rutherford Explained It n Because so few particles were deflected, they proposed that the atom has a small, dense, positively charged central core, called a nucleus. n The dense part of the atom is where most of the mass is located, and it has a positive charge.

43. The Nuclear Model of the Atom n n The new model of the atom as pictured by Rutherford’s group in 1911 is shown below. Draw Diagram in your notes Model Depicted by the Gold Foil Experiment

44. Isotopes n n In 1910, J.J. Thomson discovered that neon consisted of atoms of two different masses.

45. Isotopes Example 12 6 C 12 6 14 6 n Carbon-12 and Carbon-14 are isotopes of one another. n They are the same element with different masses. C-12C-14 Copy example in your notes

46. Isotopes n n Atoms of an element that are chemically alike but differ in mass are called isotopes of the element. n n Because of the discovery of isotopes, scientists hypothesized that atoms contained still a third type of particle that explained these differences in mass.

47. The Neutron n n Calculations showed that such a particle should have a mass equal to that of a proton but no electrical charge. n n The existence of this neutral particle, called a neutron, was confirmed in the early 1930s. n n James Chadwick is given credit for discovering the neutron.

48. The Neutron n n Isotopes of an element differ in the number of the subatomic particle called neutrons.

49. Naming Isotopes n Put the mass number as a whole number after the symbol of the element.  Carbon – 12.000000 amu  Carbon –14.003242 amu  Carbon – 14.003242 amu  Lead – 209.98418 amu  Lead – 211.99188 amu C-12 C-14 Pb-210 Pb-212

50. Subatomic Particles Electron Proton Neutron Name SymbolCharge e-e- p+p+ n0n0 +1 0 1/2000 1 1 SymbolCharge Relative mass (amu) ***REMEMBER THIS CHART***

51. Modern View of the Atom n The atom has two regions and is 3-dimensional. n The nucleus is at the center and contains the protons and neutrons.

52. Modern View of the Atom n The electron cloud is the region where you might find an electron and most of the volume of an atom.

53. Atomic Number n n The atomic number (Z) of an element is the number of protons in the nucleus of an atom of that element. n n The number of protons determines the identity of an element, as well as many of its chemical and physical properties.

54. Atomic Number n n Because atoms have no overall electrical charge, an atom must have as many electrons as there are protons in its nucleus. n n Therefore, the atomic number of an element also tells the number of electrons in a neutral atom of that element.

55. Masses n n The mass of a neutron is almost the same as the mass of a proton. n n The sum of the protons and neutrons in the nucleus is the mass number of that particular atom. n n The mass number (A) is the atomic mass rounded to a whole number.

56. SYMBOLS

57. Isotopes n Remember, isotopes of an element have different mass numbers because they have different numbers of neutrons, but they all have the same atomic number.

58. Isotopes n Subtract the atomic number from the mass number to determine the number of neutrons. n How many neutrons are in each lithium isotope below? 3 neutrons 4 neutrons 5 neutrons

59. Summary of Isotopes (copy in your notes) n # protons = atomic # n # electrons = atomic # (in neutral atoms) n # protons = # electrons (in neutral atoms) n mass # = # protons + # neutrons n # neutrons = mass # - # protons

60. Information in the Periodic Table n n The average atomic mass is the weighted average mass of all the naturally occurring isotopes of that element.

61. Average Atomic Mass n n Standard Chemistry-You are NOT responsible for knowing how to calculate average atomic mass, although a detailed example follows.

62. Calculating Atomic Mass

63. n n Copper exists as a mixture of two isotopes. n n The lighter isotope (Cu-63), with 29 protons and 34 neutrons, makes up 69.17% of copper atoms. n n The heavier isotope (Cu-65), with 29 protons and 36 neutrons, constitutes the remaining 30.83% of copper atoms.

64. Calculating Atomic Mass n n To determine the average atomic mass, first calculate the contribution of each isotope to the average atomic mass, being sure to convert each percent to a fractional abundance.

65. Calculating Atomic Mass Mass contribution = mass of isotope x abundance of isotope For Cu-63: Mass contribution = 62.930 amu x 0.6917 = 43.529 amu For Cu-65: Mass contribution = 64.928 amu x 0.3083 = 20.017 amu

66. Calculating Atomic Mass n n The average atomic mass of the element is the sum of the mass contributions of each isotope. Atomic mass Cu = mass contribution Cu-63 + mass contribution Cu-65 Atomic mass Cu = 43.529 + 20.017 = 63.546 amu

67. SYMBOLS

68. Symbols Example n Determine the complete symbol for a fluorine atom with a mass number of 19. F 19 9

69. Symbols Example n Determine the following for the fluorine atom depicted below. F 19 9 e) mass number d)atomic number c)number of electrons b)number of neutrons (9) (10) (9) (19) a)number of protons

70. Symbols Example n Determine the complete symbol for a bromine atom with a mass number of 80. Br 80 35

71. Symbols Problem n Determine the following for the bromine atom depicted below. Br 80 35 e) mass number d)atomic number c)number of electrons b)number of neutrons (35) (45) (35) (80) a)number of protons

72. Symbols Problem n If an element has an atomic number of 34 and a mass number of 78 what is the Se 78 34 d)complete symbol c)number of electrons b)number of neutrons (34) (44) (34) a)number of protons

73. Symbols Problem n If an element has 91 protons and 140 neutrons what is the Pa 231 91 d)complete symbol c)number of electrons b)mass number (91) (231) (91) a)atomic number

74. Symbols Problem n If an element has 78 electrons and 117 neutrons what is the Pt 195 78 d)complete symbol c)number of protons b)mass number (78) (195) (78) a)atomic number

75. Element# of Protons # of Neutrons # of Electrons Mass # Atomic # Potas- sium Argon n Fill in the chart below. 193919 20 184018 22

76. Element# of Protons # of Neutrons # of Electrons Mass # Atomic # Chlorine Oxygen n Fill in the chart below. (When numbers are provided, the isotope represented by each space may NOT be the most common isotope or the one closest in atomic mass to the value on the periodic table. n Fill in the chart below. (When numbers are provided, the isotope represented by each space may NOT be the most common isotope or the one closest in atomic mass to the value on the periodic table.) 173717 20 8188 8 10

77. End of Day 1