2. The Atom Chapter 4 http://www.youtube.com/watch?v =Uy0m7jnyv6U

3. I. History of the Atomic Theory A. Democritus B. Aristotle C. Lavoisier D. Proust E. Dalton F. Modern Atomic Theory II. History of Atomic Structure A. Thomson B. Milikan C. Rutherford D. Bohr E. Chadwick F. Quantum Atom III. Subatomic Particles A. Atomic Number B. Mass Number and Isotopes C. Electrons and Ions D. Nuclear and Hyphenation Notation E. Average Atomic Mass IV. Weighing and Counting Atoms A. Mole Atoms B. Mole Mass C. Mass Atoms

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 I. History of the Atomic Theory

6. B. Aristotle (384-322 BC ) & Others 1. Four kinds of matter 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) 1. 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 2. Law of Conservation of Mass a. Matter cannot be created or destroyed (in a chemical or physical change) I. History of the Atomic Theory

10. D. Joseph Proust (1779) 1. 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 – WHY???? in a whole number ratio – WHY???? I. History of the Atomic Theory

11. E. John Dalton (1803) 1. 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 2. Example: carbon monoxide carbon dioxide 2. Example: carbon monoxide 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 I. History of the Atomic Theory - Dalton

12. 3. Dalton collects data and develops his atomic theory in 1803 atomic theory in 1803 I. History of the Atomic Theory- Dalton

13. 4. Dalton’s Background a. Dalton became a school teacher at the age of 12 (he left school at age 11) b. loved meteorology - pioneer in this field c. studied works of Democritus, Boyle and Proust d. Wrote New System of Chemical Philosophy in 1808

14. 1. Matter is made of small particles-atoms 1. Matter is made of small particles-atoms 2. Atoms of a given element are identical in size, mass, but differ from those of other elements*. 3. Atoms cannot be subdivided or destroyed*. ( supports law of conservation of mass) 4.Atoms combine in small whole number ratios to form compounds. (def comp,Mult prop) 5. Atoms combine, separate, or rearrange in chemical reactions. * Modified in Modern Atomic Theory 5. Dalton’s Atomic Theory

15. JUST A THEORY……. But it lead to the But it lead to the Modern Atomic theory Modern Atomic theory

16. 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

17. 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

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

19. II.History of the Atomic Structure A. J.J. Thomson (1856-1940) J.J. Thomson (1887) 1. Experiments with cathode ray tubes a. atoms have (-) charged particles a. atoms have (-) charged particles which are smaller than atoms which are smaller than atoms b. determined charge/mass ratio of the b. determined charge/mass ratio of the “electron” “electron”

20. Voltage source +- Vacuum tube Metal Disks

21. Voltage source +-

22. +-

23. +-

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

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

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

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

28. By adding an electric field 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 + -

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

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

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

35. Demonstration of the cathode ray experiment. Demonstration of the cathode ray experiment.

36. 2.Thomson’s Model The Pudding Model 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)

37. II.History of the Atomic Structure B. Robert Milikan (1868-1953) 1. Oil Drop Experiment (1909) Oil Drop ExperimentOil Drop Experiment a. Discovered mass and actual charge of a. Discovered mass and actual charge of electron (-1) electron (-1) b. Mass is 1/1840 the mass of a hydrogen b. Mass is 1/1840 the mass of a hydrogen atom atom 1) e – has a mass of 9.11 x 10 -28 g 1) e – has a mass of 9.11 x 10 -28 g

38. Oil Drop

39. So, at this point we know: So, at this point we know: - Atoms are divisible into smaller particles - Atoms are divisible into smaller 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

40. C. Ernest Rutherford (1909) 1. Discovered the proton p + 2. Received Nobel Prize in Chemistry 3. 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 II. History of the Atomic Structure Ernest Rutherford (1871-1937)

41. Lead block Uranium Gold Foil Florescent Screen

42. He thought this would happen:

43. According to Thomson Model

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

46. Here is what he observed:

47. 4. 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

48. The positive region accounts for deflection

49. 5. 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 Rutherford Model of the Atom Rutherford Model of the Atom

50. 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 are balance! But, But, –How are the electrons arranged? –There is still mass that is unaccounted for II. History of the Atomic Structure

51. D. Niels Bohr (1913) 1. Electrons orbit nucleus in predictable paths II. History of the Atomic Structure

52. E. Chadwick (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 a. slightly greater than the mass of a proton a. slightly greater than the mass of a proton II.History of the Atomic Structure E. Chadwick (1891 – 1974)

53. F. The Quantum Atom Theory 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 II. History of the Atomic Structure

54. 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

55. 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 A. Comparing Particles 1amu

56. 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

57. 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 III.Subatomic Particles B. Atomic Number and Mass Number

58. 1. Electrons and Ions a. For neutral atoms, #e - = #p + a. For neutral atoms, #e - = #p + b. If there are more electrons, a negative ion b. If there are more electrons, a negative ion forms (anion) forms (anion) c. If there are less electrons, a positive ion c. If there are less electrons, a positive ion forms (cation) forms (cation) For now, we will work only with neutral atoms III.Subatomic Particles C. Ions

59. ions

60. Subatomic Particles C. Formation of Ions 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-

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

63. 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 III.Subatomic Particles D. Changing Number of Particles

65. 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 III.Subatomic Particles E. Nuclear Notation X Mass number Atomic number

66. How many protons? How many protons? How many neutrons? How many neutrons? How many electrons? How many electrons? III.Subatomic Particles E. Nuclear Notation Na 23 11

67. 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 III.Subatomic Particles F. Hyphen Notation

68. 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

69. 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 masses of all isotopes is added to get masses of all isotopes is added to get the average atomic mass the average atomic mass IV. Mass of Atoms B. Average Atomic Mass

70. IV. Mass of Atoms 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

71. Average Atomic Mass What is average atomic mass? Average atomic mass is the weighted average of the atomic masses of the naturally occurring isotopes of an element Average atomic mass is the weighted average of the atomic masses of the naturally occurring isotopes of an element

72. Calculating a Weighted Average Example A box contains two size of marbles. If 25.0% have masses of 2.00 g and 75.0% have masses of 3.00 g what is the weighted average? (.250) (2.00) + (.750) (3.00) =.500 + 2.25 = 2.75g 2.75g

73. Calculating Average Atomic Mass EXAMPLE Boron has two isotopes: B-10 (mass 10.013 amu) 19.8% abundance B-11 (mass 11.009 amu) 80.2% abundance Calculate the average atomic mass. (.198) (10.013) + (.802) ( 11.009) = 1.98 amu + 8.83 amu = 10.81 amu 1.98 amu + 8.83 amu = 10.81 amu

74. Calculating Average Atomic Mass Calculate the average atomic mass of Mg. Isotope 1 - 23.985 amu (78.99%) Isotope 2 - 24.986 amu (10.00%) Isotope 3 – 25.982 amu (11.01%) (23.985)(.7899)+(24.986)(.1000)+(25.982)(.1101) 18.95 amu + 2.498 amu + 2.861 amu = 24.31 amu

75. Average Atomic Mass Helium has two naturally occurring isotopes, He-3 and He-4. The atomic mass of helium is 4.003 amu. Which isotope is more abundant in nature? He-4 is more abundant in nature because the atomic mass is closer to the mass of He-4 than to the mass of He-3.

76. Calculating Average Atomic Mass Process 1. Multiply %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

77. Isotopic Pennies – number of pre and post 1982 Isotopic Pennies – number of pre and post 1982 a. Let X be the number of pre-1982 pennies b. Let 10-X be the number of post-1982 pennies c. (X)(3.1g) + (10-X)(2.5g) = mass of 10 pennies pre-82 post-82 pre-82 post-82 EXAMPLE (Mass of a sample of pennies is 31.0g) (X)(3.1g) + [(10-X)(2.5g)] = 31.0 g 3.1X + 25 - 2.5X = 31.0g.6X + 25 = 31.0g.6X = 6.0g X = 6.0g/.6 X = 10 pre-82 pennies 10-X = 0 post-82 pennies

78. Isotopic Penny Lab- Average Atomic Mass Calculate percent of pre-82 and post-82 pennies # of pre-82 pennies x 100% # post-82 pennies x 100% 10 10 10 10 Calculate the average atomic mass of coinium (% pre-82)(3.1g) + (% post-82)(2.5) = average atomic mass

79. V.Radioactive Decay A. The Process 1. What is radioactive decay? a. spontaneous nuclear change in which unstable nuclei emit radiation and lose energy a. spontaneous nuclear change in which unstable nuclei emit radiation and lose energy 1) radiation – rays and particles 1) radiation – rays and particles emitted by radioactive materials emitted by radioactive materials 2. Why do atoms undergo decay? a. produce a nucleus that is more stable a. produce a nucleus that is more stable

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

81. Alpha Particle

82. 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

83. 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)

84. Beta Decay

86. + beta particle

87. 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)

88. Alpha Decay

91. + alpha particle

92. Radioactive Decay of Uranium

94. Alpha Decay + alpha particle

95. Radiation

96. Radiation

97. Radiation

98. The Mole What is a mole in chemistry? What is a mole in chemistry? What conversion factors are associated with the mole? What conversion factors are associated with the mole? Types of conversions involving mole equalities Types of conversions involving mole equalities

99. What is a Mole? What are mole equalities A mole is equal to 6.02 x 10 23 particles A mole is equal to 6.02 x 10 23 particles Particles can be atoms, molecules or ions Particles can be atoms, molecules or ions 6.02 x 10 23 is Avogadro's Number 6.02 x 10 23 is Avogadro's Number Mole Equalities Mole Equalities - 1 mole = molar mass - 1 mole = molar mass - 1 mole = 6.02 x 10 23 particles - 1 mole = 6.02 x 10 23 particles

100. 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)

101. 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

102. Solving Mole Problems EXAMPLES 1. 1.00 mole of He = 4.00 g. 2. 2.00 mole of He = _____g 2.00 mol He X 4.00g He = 2.00 mol He X 4.00g He = 1 1 mole He 1 1 mole He 8.00 g He 8.00 g He

103. EXAMPLES 3. 1.00 mole He = 6.02 X 10 23 atoms 4. 2.00 mole He = ________atoms He 2.00 mole He x 6.02 x 10 23 atoms He = 12.04 x 10 23 1 1 mole He 1 1 mole He 1.20 x 10 24 atoms He 1.20 x 10 24 atoms He 5. 16.00g He = _____ moles He 16.0 g He x 1 mole He = 1 4.00 g He 4.00 moles He 1 4.00 g He 4.00 moles He

104. EXAMPLES 6. 3.01 X 10 23 atoms He = _____ moles 3.01 x 10 23 atoms He x 1 mole He = 1 6.02 x 10 23 atoms He 1 6.02 x 10 23 atoms He.500 mol He.500 mol He 7. 8.00g He =______atoms He 8.00 g He x 1 mole He x 6.02 x 10 23 atoms He = 1 4.00g He 1 mole He 1 4.00g He 1 mole He 12.04 x 10 23 atoms He = 1.20 x 10 24 atoms He 12.04 x 10 23 atoms He = 1.20 x 10 24 atoms He

105. Sample Problems 1. Moles to mass. Find the mass of 3.50 moles of carbon. Find the mass of 3.50 moles of carbon. 2. Mass to Moles How many moles of carbon are contained in How many moles of carbon are contained in 60.0 g of carbon? 60.0 g of carbon? 3. Moles to Atoms How many atoms of carbon are found in 4.00 moles of carbon? How many atoms of carbon are found in 4.00 moles of carbon?

106. Sample Problems 4. Atoms to Moles How many moles of carbon are represented by How many moles of carbon are represented by 1.806 x 10 24 atoms of carbon? 1.806 x 10 24 atoms of carbon? 5. Mass to Atoms How many carbon atoms are found in 36.0g of How many carbon atoms are found in 36.0g of carbon? carbon? 6. Atoms to Mass What is the mass of 1.204 x 10 24 atoms of carbon? What is the mass of 1.204 x 10 24 atoms of carbon?

107. More Sample Problems 2.00 moles of Cu = atoms of Cu 60.0 grams of C = moles of C 3.00 x 10 23 atoms He = moles of He 2.50 moles Al = grams of Al 28.0 grams N = atoms of N 1.80 x 10 23 atoms Mg = grams of Mg

108. 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.