2. The Atom Chapter 4

3. I. History of the Atomic Theory IV. Mass of Atoms A. Democritus A. Atomic Mass B. Aristotle B. Average Atomic Mass C. Lavoisier C. Mole and Molar Mass D. Proust E. Dalton F. Modern Atomic Theory II. History of Atomic Structure V. Radioactive Decay A. Thomson A. Process B. Millikan B. Comparison C. Rutherford C. Types D. Bohr E. Chadwick F. Quantum Atom III. Subatomic Particles A. Comparing Particles B. Atomic Number and Mass Number C. Ions D. Changing Number of Particles E. Nuclear Notation E. Nuclear Notation F. Hyphen Notation F. Hyphen Notation

4. I. History of the Atomic Theory Remember: a scientific theory explains behaviors and the ‘nature’ of things Remember: a scientific theory explains behaviors and the ‘nature’ of things Theories can be revised when new discoveries are made Theories can be revised when new discoveries are made The theory describing the composition of matter has been revised many times The theory describing the composition of matter has been revised many times

5. A. Democritus (460-370 BC ) 1.Matter is made up of “atoms” that are solid, indivisible and indestructible 2.Atoms constantly move in space 3.Different atoms have different size and shape 4.Changes in matter result from changes in the grouping of atoms 5. Properties of matter result from size, shape and movement

6. B.Aristotle (384-322 BC ) 1. Four kinds of matter a. Fire – Earth – Water – Air a. Fire – Earth – Water – Air 2. One kind of matter can transform 2. One kind of matter can transform into another into another 3. Rejected idea of the “atom” (idea then 3. Rejected idea of the “atom” (idea then ignored for almost 2000 years ignored for almost 2000 years 4. This theory was more popular and 4. This theory was more popular and it was easier to accept it was easier to accept

7. Aristotle’s Theory of Matter

8. C. Antoine Lavoisier (1770s) Experiment: Experiment: 2 Sn + O 2  2 SnO 2 Sn + O 2  2 SnO tin oxygen tin (II) oxide tin oxygen tin (II) oxide mass before reaction = mass after reaction Law of Conservation of Mass Matter cannot be created or destroyed (in a chemical or physical change)

9. 1788-1799

10. D. Joseph Proust (1779) Develops Law of Definite Composition- all samples of a specific substance contain the same mass ratio of the same elements a. ex: all samples of CO 2 contains 27.3% a. ex: all samples of CO 2 contains 27.3% carbon and 72.7% oxygen carbon and 72.7% oxygen b. therefore ‘elements’ are combining b. therefore ‘elements’ are combining in a whole number ratio in a whole number ratio

11. D. John Dalton (1803) Dalton became a school teacher at the age of 12 (he left school at age 11) Loved meteorology - pioneer in this field Studied works of Democritus, Boyle and Proust Wrote New System of Chemical Philosophy in 1808

12. Develops Law of Multiple Proportions a. describes the ratio of elements by mass in two different compounds composed of the same elements a. describes the ratio of elements by mass in two different compounds composed of the same elements Example: Example: carbon monoxide vs. carbon dioxide carbon monoxide vs. carbon dioxide 1 part oxygen : 2 parts oxygen 1 part oxygen : 2 parts oxygen *when compared to the same amount of carbon in each compound carbon in each compound

13. a. Matter is made of small particles-atoms b. Atoms of a given element are identical in size, mass, but differ from those of other elements*. c. Atoms cannot be subdivided or destroyed*. d. Atoms combine in small whole number ratios to form compounds. e. Atoms combine, separate, or rearrange in chemical reactions. * Modified in Modern Atomic Theory * Modified in Modern Atomic Theory Dalton collects data and develops Atomic Theory in 1803

14. 1.All matter is made up of small particles called atoms. 2.Atoms of the same element have the same chemical properties while atoms of different elements have different properties 3.Not all atoms of an element have the same mass, but they all have a definite average mass which is characteristic. (isotopes) F. Modern Atomic Theory (1)

15. 4. Atoms of different elements combine to form compounds and each element in the compound loses its characteristic properties. 5. Atoms cannot be subdivided by chemical or physical changes – only by nuclear changes changes F. Modern Atomic Theory (2)

16. II.History of the Atomic Structure A. J.J. Thomson (1887) Experiments with cathode ray tubes

17. Voltage source +- Vacuum tube Metal Disks

18. Voltage source +-

19. +-

20. +-

21. n Passing an electric current makes a beam appear to move from the negative to the positive end Voltage source +-

22. n Passing an electric current makes a beam appear to move from the negative to the positive end Voltage source +-

23. n Passing an electric current makes a beam appear to move from the negative to the positive end Voltage source +-

24. n Passing an electric current makes a beam appear to move from the negative to the positive end Voltage source +-

25. By adding an electric field By adding an electric field

26. Voltage source n By adding an electric field + -

27. Voltage source n By adding an electric field + -

28. Voltage source n By adding an electric field + -

29. Voltage source n By adding an electric field + -

30. Voltage source n By adding an electric field + -

31. Voltage source n By adding an electric field he found that the moving pieces were negative + -

32. Thomson’s Model of the Atom a. electrons present (-) a. electrons present (-) b. atom is like plum pudding - bunch of positive stuff (pudding), with the electrons suspended (plums) b. atom is like plum pudding - bunch of positive stuff (pudding), with the electrons suspended (plums) Calculated the ratio between the charge of the electron and its mass: e/m Calculated the ratio between the charge of the electron and its mass: e/m “Chocolate Chip Cookie” or “Plum Pudding Model”

33. Millikan’s Oil Drop Experiments Robert Milikan (1909) –Oil Drop Experiment Oil Drop ExperimentOil Drop Experiment –Measured the electrical charge on the electron –Mass can be calculated (Thomson determined the e/m ratio) –Mass is 1/1840 the mass of a hydrogen atom –electron has a mass of 9.11 x 10 -28 g

34. B. Millikan’s Oil Drop Experiment (1909)

35. Oil-Drop Experiment: Animation http://wps.prenhall.com/wps/media/object s/602/616761/MillikanOilDropExperiment.h tml http://wps.prenhall.com/wps/media/object s/602/616761/MillikanOilDropExperiment.h tml http://wps.prenhall.com/wps/media/object s/602/616761/MillikanOilDropExperiment.h tml http://wps.prenhall.com/wps/media/object s/602/616761/MillikanOilDropExperiment.h tml http://www.britannica.com/nobel/cap/omil lik001a4.html http://www.britannica.com/nobel/cap/omil lik001a4.html http://www.britannica.com/nobel/cap/omil lik001a4.html http://www.britannica.com/nobel/cap/omil lik001a4.html http://www.daedalon.com/oildrop.html http://www.daedalon.com/oildrop.html http://www.daedalon.com/oildrop.html

36. So, at this point we know: So, at this point we know: - Atoms are indivisible particles - Atoms are indivisible particles –Electrons are negatively charged –The mass of an electron is very small HOWEVER HOWEVER –Atoms should have a (+) portion to balance the negative part the negative part - Electrons are so small that some other particles must account for mass II.History of the Atomic Structure – Summary thus far

37. Experiment (1909) Experiment (1909) Gold Foil Experiment (Expectations) Gold Foil Experiment (Expectations)Gold Foil ExperimentGold Foil Experiment a. Shot alpha particles at atoms of gold a. Shot alpha particles at atoms of gold b. expected them to pass straight b. expected them to pass straight through through C. Ernest Rutherford (1871-1937)

38. Lead block Uranium Gold Foil Florescent Screen

39. He thought this would happen:

40. According to Thomson Model

41. He thought the mass of the positive charge was evenly distributed in the atom

43. Here is what he observed:

44. The positive region accounts for deflection

45. Gold Foil Experiment Results a. Most positive alpha particles pass right through through b. However, a few were deflected c. Rutherford reasoned that the positive alpha particle was deflected or repelled alpha particle was deflected or repelled by a concentration of positive charge by a concentration of positive charge

46. Gold Foil Experiment Conclusions a. the atom is mostly empty space a. the atom is mostly empty space b. the atom has a small, dense positive center b. the atom has a small, dense positive center surrounded by electrons surrounded by electrons

47. At this point in 1909, we know: At this point in 1909, we know: –p + = 1.67 x 10 -24 g –e - = 9.11 x 10 -28 g –The charges balance! But, But, –How are the electrons arranged? –There is still mass that is unaccounted for II. History of the Atomic Structure

48. Rutherford Model of the Atom

49. Electrons orbit nucleus in predictable paths D. Niels Bohr (1913)

50. In 1935 In 1935 1. Discovers neutron in nucleus 2. Neutron is neutral - does not have a charge n 0 not have a charge n 0 3. Mass is 1.67 x 10 -24 g slightly greater than the mass of a proton slightly greater than the mass of a proton E. Chadwick (1891 – 1974)

51. History of the Atomic Theory History of the Atomic Theory 18031897190919131935Today solidparticleelectronproton e- orbit nucleus neutron Quantum Atom theory DaltonThomsonRutherfordBohrChadwick Schrodinger and others

52. Charges balanced Charges balanced Mass accounted for Mass accounted for However – However – what about the what about the behavior of the behavior of the electrons? electrons? II. History of the Atomic Structure

53. 1. The atom is mostly empty space space 2. Two regions: a. Nucleus- protons and neutrons b. Electron cloud- region where you have a 90% chance of finding an electron F. The Quantum Atom Theory

54. Electron Proton Neutron NameSymbolCharge Relative mass Actual mass (g) e-e- p+p+ n0n0 +1 0 0 1amu 9.11 x 10 -28 1.67 x 10 -24 III.Subatomic Particles 1amu Comparing Particles

55. 1. Atomic number 1. the number of protons in the nucleus of an atom a. identifies the element a. identifies the element b. no two elements have the same atomic number b. no two elements have the same atomic number 2.Ex. C is 6, N is 7 and O is 8 carbon nitrogen oxygen carbon nitrogen oxygen III.Subatomic Particles B. Atomic Number and Mass Number

56. 2. Mass number a. the number of protons plus neutrons in the nucleus of an atom b. mass number is very close to the mass of an atom in amu (atomic mass units) c. two atoms with the same atomic number but different mass number are called isotopes 1) (mass #) – (atomic #) = #n 0 1) (mass #) – (atomic #) = #n 0 B. Atomic Number and Mass Number

57. 1. Electrons and Ions a. For neutral atoms, #e - = #p + a. For neutral atoms, #e - = #p + b. When there are more electrons than b. When there are more electrons than protons, a negative ion forms (anion) protons, a negative ion forms (anion) c. When there are less electrons than protons, c. When there are less electrons than protons, a positive ion forms (cation) a positive ion forms (cation) For now, we will work only with neutral atoms C. Formation of Ions

58. Examples of Ions Atom loses electrons Atom gain electrons and form cations and forms anions Cations (+ ions) Anions (- ions) K + Br - K + Br - Ca 2+ O 2- Ca 2+ O 2- Al 3+ N 3- Al 3+ N 3-

59. C. Formation of Ions From Atoms Na loses an electron and forms a cation Na 0 – e- --> Na+ Cl gains an electron and forms an anion Cl 0 + e- --> Cl-

60. 1. You can never change the number of protons and have the same element protons and have the same element 2. If you change the number of neutrons in an atom, you get an isotope 3. If you change the number of electrons in an atom, you get an ion D. Changing Number of Particles

61. 1.Nuclear Notation is one method for depicting isotopes of an element 2.contains the symbol of the element, the mass number, and the atomic number E. Nuclear Notation X Mass number Atomic number

62. How many protons? How many protons? How many neutrons? How many neutrons? How many electrons? How many electrons? Na 23 11

63. 1. Element symbol or name – mass # 2. EXAMPLES a. Fluorine-19 a. Fluorine-19 b. C-14 b. C-14 c. U-238 c. U-238 F. Hyphen Notation

64. IV.Mass of Atoms A. Atomic Mass 1. Mass of an atom a. too small to measure in grams a. too small to measure in grams b. use relative mass (amu) b. use relative mass (amu) 1) atomic mass unit 1) atomic mass unit 2) 1 amu is defined as 1/12 the mass 2) 1 amu is defined as 1/12 the mass of one C-12 atom of one C-12 atom

65. 1. weighted average mass of all known isotopes a. weighted means that the frequency of an isotope is considered a. weighted means that the frequency of an isotope is considered b. mass of each isotope is multiplied by its percent occurrence in nature – then its percent occurrence in nature – then product of all isotopes is added to get product of all isotopes is added to get the average atomic mass the average atomic mass B. Average Atomic Mass

66. Calculating Average Atomic Mass Process 1. Mulitply % occurrence x mass of isotope 2. Add products for each isotope isotope occurrence isotope occurrence isotope occurrence isotope occurrence Ex. X- 40 (30.0% ) X-30 (70.0%) (40 x.300) + (30 x.700) = (40 x.300) + (30 x.700) = 12 + 21 = 33 amu 12 + 21 = 33 amu

67. C. The Mole and Molar Mass 1. measures the amount of substance a. 1 mole = 6.02x10 23 (Avogdro’s #) of a. 1 mole = 6.02x10 23 (Avogdro’s #) of particles (atoms, molecules, ions, electrons) particles (atoms, molecules, ions, electrons) b. standard – 1mole is the number of atoms in b. standard – 1mole is the number of atoms in 12g of C-12 isotope 12g of C-12 isotope 2. Molar mass – mass in grams of one mole (mol) of any substance (mol) of any substance a. numerically equal to atomic mass in amu a. numerically equal to atomic mass in amu b. unit is grams/mol b. unit is grams/mol

68. V.Radioactive Decay A. The Process What is radioactive decay? spontaneous nuclear change in which unstable nuclei emit radiation and lose energy spontaneous nuclear change in which unstable nuclei emit radiation and lose energy radiation – rays and particles radiation – rays and particles emitted by radioactive materials emitted by radioactive materials Why do atoms undergo decay? Why do atoms undergo decay? produce a nucleus that is more stable- p/n ratio produce a nucleus that is more stable- p/n ratio

69. B. Comparison of alpha, beta, gamma Alpha Beta Gamma____ Alpha Beta Gamma____ Formparticleparticle electromagnetic radiation radiationSymbol Mass4 amuno massno mass Charge+2-1none Notation

70. Alpha Particle

71. V.Radioactive Decay C. Types of Decay 1. Beta Decay (neutron  proton + electron) a. beta particle (electron) is given off a. beta particle (electron) is given off b. atomic number increases by one b. atomic number increases by one c. mass number stays the same c. mass number stays the same 2. Alpha Decay a. alpha particle (2p + +2n 0 ) is given off a. alpha particle (2p + +2n 0 ) is given off b. atomic number decreases by 2 b. atomic number decreases by 2 c. mass number decreases by 4 c. mass number decreases by 4

72. V.Radioactive Decay D. Examples of Decay Beta Decay (n 0  p + + e - ) e - released Parent Nuclei Daughter Nuclei Co-60 --------------> Ni-60 + e - Co-60 --------------> Ni-60 + e - (z = 27) (z = 28) (z = 27) (z = 28) C-14 ---------------> N-14 + e - C-14 ---------------> N-14 + e - (z = 6) (z = 7) (z = 6) (z = 7)

73. Beta Decay

75. + beta particle

76. Radioactive Decay D. Examples of Decay Alpha Decay (alpha particle (2n 0 + 2p + ) released) released) Parent Nuclei Daughter Nuclei Th-232 ----------------> Ra-228 + alpha Th-232 ----------------> Ra-228 + alpha (z = 90) (z = 88) (z = 90) (z = 88) Ra-226 ---------------> Rn-222 + alpha Ra-226 ---------------> Rn-222 + alpha (z = 88) (z = 86) (z = 88) (z = 86)

77. Alpha Decay

80. + alpha particle

81. Radioactive Decay of Uranium

83. Radiation

84. Radiation

85. Radiation

86. Mole Calculations Using Conversion Factors 1 mole molar mass 1 mole molar mass 6.02 x 10 23 1 mole 6.02 x 10 23 1 mole PARTICLES MOLES MASS 6.02 x 10 23 1 mole 6.02 x 10 23 1 mole 1 mole molar mass 1 mole molar mass

87. Mole Conversions [mass-mole-atoms] Type Equality Used 1. MOLES  MASS 2. MASS  MOLES 1 mole= molar mass (g) 3. MOLES  ATOMS 4. ATOMS  MOLES 1 mole = 6.02 x 10 23 atoms 5. MASS  ATOMS 1 mole = 6.02 x 10 23 atoms 6. ATOMS  MASS 1 mole = molar mass (g)

88. Mole Calculations (Mass of Helium is 4.00 ) 1. 1.00 mole of helium = 4.00g 2. 2.00 mole of helium = 8.00g 3. 1.00 mole of helium = 6.02 x 10 23 atoms 4. 2.00 mole of helium = 1.20 x 10 24 atoms 5. 16.0g of helium = 4.00 mol 6. 3.0l x 10 23 atoms of helium =.500 moles 7. 8.00g of helium = 1.20 x 10 24 atoms.