1. Early Atomic Theory and Structure Chapter 5
Hein and Arena
Eugene Passer Chemistry Department Bronx Community College
© John Wiley and Sons, Inc
Version 2.0
12th Edition

2. Chapter Outline 5.1 Early Thoughts 5.6 Subatomic Parts of the Atom
5.2 Dalton's Model of the Atom
5.7 The Nuclear Atom
5.3 Composition of Compounds
5.8 Isotopes of the Elements
5.4 The Nature of Electric Charge
5.9 Atomic Mass
5.5 Discovery of Ions

3. 5.1 Early Thoughts

4. The earliest models of the atom were developed by the ancient Greek philosophers.
Empedocles stated that matter was made of 4 elements: earth, air, fire, and water.
Democritus (about B.C.) thought that all forms of matter were divisible into tiny indivisible particles. He called them “atoms” from the Greek “atomos”, indivisible.

5. Aristotle (384-322 B.C.) rejected the theory of Democritus and advanced the Empedoclean theory.
Aristotle’s influence dominated the thinking of scientists and philosophers until the beginning of the 17th century.

6. 5.2 Dalton’s Model of the Atom

7. 2000 years after Aristotle, John Dalton, an English schoolmaster, proposed his model of the atom–which was based on experimentation.

8. Dalton’s Atomic Theory
Elements are composed of minute indivisible particles called atoms.
Atoms under special circumstances can be decomposed.
Modern research has demonstrated that atoms are composed of subatomic particles.
Atoms of the same element are alike in mass and size.
Atoms of different elements have different masses and sizes.
Chemical compounds are formed by the union of two or atoms of different elements.

9. Dalton’s Atomic Theory
Atoms combine to form compounds in simple numerical ratios, such as one to one , two to two, two to three, and so on.
Atoms of two elements may combine in different ratios to form more than one compound.

10. Dalton’s atoms were individual particles.
Atoms of each element are alike in mass and size.
5.1

11. Dalton’s atoms were individual particles.
Atoms of different elements are not alike in mass and size.
5.1

12. Dalton’s atoms combine in specific ratios to form compounds.

13. 5.3 Composition of Compounds

14. The Law of Definite Composition
A compound always contains two or more elements chemically combined in a definite proportion by mass.

15. Composition of Water
Water always contains the same two elements: hydrogen and oxygen.
The percent by mass of hydrogen in water is 11.2%.
The percent by mass of oxygen in water is 88.8%.
Water always has these percentages. If the percentages were different, the compound would not be water.

16. Composition of Hydrogen Peroxide
Hydrogen peroxide always contains the same two elements: hydrogen and oxygen.
The percent by mass of hydrogen in hydrogen peroxide is 5.9%.
The percent by mass of oxygen in hydrogen peroxide is 94.1%.
Hydrogen peroxide always has these percentages. If the percentages were different, the compound would not be hydrogen peroxide.

17. The Law of Multiple Proportions
Atoms of two or more elements may combine in different ratios to produce more than one compound.

18. Combining Masses of Hydrogen and Oxygen
Mass Hydrogen(g)
Mass Oxygen(g)
Water
1.0
8.0
Hydrogen Peroxide
16.0
Hydrogen peroxide has twice as much oxygen (by mass) as does water.

19. Combining Ratios of Hydrogen and Oxygen
Hydrogen peroxide has twice as many oxygens per hydrogen atom as does water.
The formula for water is H2O.
The formula for hydrogen peroxide is H2O2.

21. 5.4 The Nature of Electric Charge

22. Properties of Electric Charge
Charge may be of two types: positive and negative.
Unlike charges attract (positive attracts negative), and like charges repel (negative repels negative and positive repels positive).
Charge may be transferred from one object to another, by contact or induction.
The smaller the distance between two charges, the greater the force of attraction between unlike charges (or repulsion between identical charges).
q1 and q2 are charges, r is the distance between charges, and k is a constant.

23. 5.5 Discovery of Ions

24. Michael Faraday discovered that certain substances, when dissolved in water, conducted an electric current.
He found that atoms of some elements moved to the cathode (negative electrode) and some moved to the anode (positive electrode).
He concluded they were electrically charged and called them ions (Greek wanderer).

25. Svante Arrhenius reasoned that an ion is an atom (or a group of atoms) carrying a positive or negative electric charge.
Arrhenius accounted for the electrical conduction of molten sodium chloride (NaCl) by proposing that melted NaCl dissociated into the charged ions Na+ and Cl-.
NaCl → Na+ + Cl- Δ

26. NaCl → Na+ + Cl-
When melted, the positive Na+ ions moved to the cathode (negative electrode). Thus positive ions are called cations.
When melted, the negative Cl- ions moved to the anode (positive electrode). Thus negative ions are called anions.

27. 5.6 Subatomic Parts of the Atom

28. An atom is very small

29. The diameter of an atom is 0.1 to 0.5 nm.
This is 1 to 5 ten billionths of a meter.
Even smaller particles than atoms exist. These are called subatomic particles.
If the diameter of this dot is 1 mm, then 10 million hydrogen atoms would form a line across the dot.

30. Subatomic Particles

31. Electron

32. In 1875 Sir William Crookes invented the Crookes tube.
Crookes tubes experiments led the way to an understanding of the subatomic structure of the atom.
Crookes tube emissions are called cathode rays.

33. In 1897 Sir Joseph Thompson demonstrated that cathode rays:
travel in straight lines.
are negative in charge.
are deflected by electric and magnetic fields.
produce sharp shadows
are capable of moving a small paddle wheel.

34. This was the discovery of the fundamental unit of charge – the electron.

35. Proton

36. Eugen Goldstein, a German physicist, first observed protons in 1886:
Thompson determined the proton’s characteristics.
Thompson showed that atoms contained both positive and negative charges.
This disproved the Dalton model of the atom which held that atoms were indivisible.

37. Neutron

38. James Chadwick discovered the neutron in 1932.
Its actual mass is slightly greater than the mass of a proton.

40. Ions

41. Positive ions were explained by assuming that a neutral atom loses electrons.
Negative ions were explained by assuming that additional electrons can be added to atoms.

42. When one or more electrons are lost from an atom, a cation is formed.
5.4

43. When one or more electrons are added to a neutral atom, an anion is formed.
5.4

44. 5.7 The Nuclear Atom

45. Radioactivity was discovered by Becquerel in 1896.
Radioactive elements spontaneously emit alpha particles, beta particles and gamma rays from their nuclei.
By 1907 Rutherford found that alpha particles emitted by certain radioactive elements were helium nuclei.

46. The Rutherford Experiment

47. Rutherford in 1911 performed experiments that shot a stream of alpha particles at a gold foil.
Most of the alpha particles passed through the foil with little or no deflection.
He found that a few were deflected at large angles and some alpha particles even bounced back.

48. Rutherford’s alpha particle scattering experiment.
5.5

49. An electron with a mass of 1/1837 amu could not have deflected an alpha particle with a mass of 4 amu.
Rutherford knew that like charges repel.
Rutherford concluded that each gold atom contained a positively charged mass that occupied a tiny volume. He called this mass the nucleus.

50. If a positive alpha particle approached close enough to the positive mass it was deflected.
Most of the alpha particles passed through the gold foil. This led Rutherford to conclude that a gold atom was mostly empty space.

51. Because alpha particles have relatively high masses, the extent of the deflections led Rutherford to conclude that the nucleus was very heavy and dense.

52. Deflection Scattering
Deflection and scattering of alpha particles by positive gold nuclei.
5.5

53. General Arrangement of Subatomic Particles

54. Rutherford’s experiment showed that an atom had a dense, positively charged nucleus.
Chadwick’s work in 1932 demonstrated that the atom contains neutrons.
Rutherford also noted that light, negatively charged electrons were present in an atom and offset the positive nuclear charge.

55. Rutherford put forward a model of the atom in which a dense, positively charged nucleus is located at the atom’s center.
The negative electrons surround the nucleus.
The nucleus contains protons and neutrons.

56. 5.6

57. Atomic Numbers of the Elements

58. The atomic number of an element is equal to the number of protons in the nucleus of that element.
The atomic number of an atom determines which element the atom is.

59. Every atom with an atomic number of 1 is a hydrogen atom
Every atom with an atomic number of 1 is a hydrogen atom. Every hydrogen atom contains 1 proton in its nucleus.

60. atomic number 1H
Every atom with an atomic number of 1 is a hydrogen atom.
1 proton in the nucleus

61. Every atom with an atomic number of 6 is a carbon atom
Every atom with an atomic number of 6 is a carbon atom. Every carbon atom contains 6 protons in its nucleus.

62. atomic number 6C
Every atom with an atomic number of 6 is a carbon atom.
6 protons in the nucleus

63. atomic number
92U
Every atom with an atomic number of 92 is a uranium atom.
92 protons in the nucleus

64. 5.8 Isotopes of the Elements

65. Atoms of the same element have the same number of protons.
Atoms of the same element can have different masses, because they can have different numbers of neutrons.
These are isotopes of the same element.

66. Isotopes of the Same Element Have
Equal numbers of protons
Different numbers of neutrons

67. Isotopic Notation

68. Isotopic Notation
6 protons + 6 neutrons C 6 12
6 protons

69. Isotopic Notation
6 protons + 8 neutrons 14 C 6
6 protons

70. Isotopic Notation
8 protons + 8 neutrons 16 O 8
8 protons

71. Isotopic Notation
8 protons + 9 neutrons 17 O 8
8 protons

72. Isotopic Notation
8 protons + 10 neutrons 18 O 8
8 protons

73. Hydrogen has three isotopes
1 proton
0 neutrons
1 proton
1 neutron
1 proton
2 neutrons

74. Examples of Isotopes Element Protons Electrons Neutrons Symbol
Hydrogen H
Hydrogen H
Hydrogen H
Uranium U
Uranium U
Chlorine Cl
Chlorine Cl

75. 5.9 Atomic Mass

76. The mass of a single atom is too small to measure on a balance.
Using a mass spectrometer, the mass of the hydrogen atom was determined.

77. A Modern Mass Spectrometer
Positive ions formed from sample.
Electrical field at slits accelerates positive ions.
Deflection of positive ions occurs at magnetic field.
From the intensity and positions of the lines on the mass spectrogram, the different isotopes and their relative amounts can be determined.
A mass spectrogram is recorded.
5.8

78. A typical reading from a mass spectrometer
A typical reading from a mass spectrometer. The two principal isotopes of copper are shown with the abundance (%) given.
5.9

79. Using a mass spectrometer, the mass of one hydrogen atom was determined to be 1.673 x 10-24 g.

80. This number is very small.

81. To overcome this problem, a system of relative atomic masses using “atomic mass units” was devised to express the masses of elements using simple numbers.
Numbers of this size are too small for practical use.

82. The standard to which the masses of all other atoms are compared to was chosen to be the most abundant isotope of carbon.

83. A mass of exactly 12 atomic mass units (amu) was assigned to

84. 1 amu is defined as exactly equal to the mass of a carbon-12 atom
1 amu = x g

85. Average atomic mass 1.00797 amu.

86. Average atomic mass 39.098 amu.

87. Average atomic mass 248.029 amu.

88. Average Relative Atomic Mass

89. Most elements occur as mixtures of isotopes.
Isotopes of the same element have different masses.
The listed atomic mass of an element is the average relative mass of the isotopes of that element compared to the mass of carbon-12 (exactly …amu).

90. Average atomic mass (amu)
To calculate the atomic mass, multiply the atomic mass of each isotope by its percent abundance and add the results.
Isotope
Isotopic mass (amu)
Abundance (%)
Average atomic mass (amu)
69.09
30.91
63.55
( amu) =
43.48 amu
( amu) =
20.07 amu
63.55 amu

91. Relationship Between Mass Number and Atomic Number

92. The mass number minus the atomic number equals the number of neutrons in the nucleus.- =
109 - 47 = 62

93. The End