2. John Dalton (1766- 1844), an English schoolteacher and chemist. Dalton’s Atomic Theory Development of the Modern Atomic Theory Chapter 5: Electrons in Atoms

3. The following statements are the main points of Dalton’s atomic theory. Dalton’s Atomic Theory 1. All matter is made up of atoms. 2. Atoms are indestructible and cannot be divided into smaller particles. (Atoms are indivisible.) 3. All atoms of one element are exactly alike, but are different from atoms of other elements.

4. Dalton’s atomic theory, led most scientists in the 1800s believed that the atom was like a tiny solid ball that could not be broken up into parts. The Electron In 1897, a British physicist, J.J. Thomson, discovered that this solid-ball model was not accurate. Thomson’s experiments used a vacuum tube.

5. A vacuum tube has had all gases pumped out of it. The Electron 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. These electrodes become electrically charged when they are connected to a high-voltage electrical source.

6. When the electrodes are charged, rays travel in the tube from the negative electrode, which is the cathode, to the positive electrode, the anode. Cathode-Ray Tube These rays originate at the cathode, they are called cathode rays. Thomson found that the rays bent toward a positively charged plate and away from a negatively charged plate.

7. These electrons had to come from the matter (atoms) of the negative electrode. Cathode-Ray Tube Thomson concluded that cathode rays are made up of invisible, negatively charged particles referred to as electrons. He knew that objects with like charges repel each other, and objects with unlike charges attract each other.

8. From Thomson’s experiments, scientists had to conclude that atoms were not just neutral spheres, but somehow were composed of electrically charged particles. Cathode-Ray Tube Reason should tell you that there must be a lot more to the atom than electrons. Matter is not negatively charged, so atoms can’t be negatively charged either. If atoms contained extremely light, negatively charged particles, then they must also contain positively charged particles—probably with a much greater mass than electrons.

9. Rutherford’s Gold Foil Experiment 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 cookie-dough model of the atom.

10. In 1886, scientists discovered that a cathode- ray tube emitted rays not only from the cathode but also from the positively charged anode. Protons These rays travel in a direction opposite to that of cathode rays.

11. Like cathode rays, they are deflected by electrical and magnetic fields, but in directions opposite to the way cathode rays are deflected. Protons Thomson was able to show that these rays had a positive electrical charge. Years later, scientists determined that the rays were composed of positively charged subatomic particles called protons.

12. At this point, it seemed that atoms were made up of equal numbers of electrons and protons. Protons However, in 1910, Thomson discovered that neon consisted of atoms of two different masses. This led to isotopes. He didn’t know where they were in the atom. protons.

13. Protons Today, chemists know that neon consists of three naturally occurring isotopes. Atoms of an element that are chemically alike but differ in mass are called isotopes of the element. The third was too scarce for Thomson to detect.

14. Rutherford’s Gold Foil Experiment The experimenters set up a lead-shielded box containing radioactive polonium, which emitted a beam of positively charged subatomic particles through a small hole. 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. Today, we know that the particles of the beam consisted of clusters containing two protons and two neutrons and are called alpha particles.

15. The Gold Foil Experiment

16. The Nuclear Model of the Atom The new model of the atom as pictured by Rutherford’s group in 1911 is shown below.

17. The Nuclear Model of the Atom Most of the particles passed through the foil, they concluded that the atom is nearly all empty space. To explain the results of the experiment, Rutherford’s team proposed a new model of the atom. So few particles were deflected, they proposed that the atom has a small, dense, positively charged central core, called a nucleus.

18. Neutrons Calculations showed that such a particle should have a mass equal to that of a proton but no electrical charge. The discovery of isotopes, scientists hypothesized that atoms contained still a third type of particle that explained these differences in mass. The existence of this neutral particle, called a neutron, was confirmed in the early 1930s.

19. The Nuclear Model of the Atom Because so few particles were deflected, they proposed that the atom has a small, dense, positively charged central core, called a nucleus. Thompson 1906 Rutherford 1913 Bohr 1924

20. Atomic Numbers It is the number of protons that determines the identity of an element, as well as many of its chemical and physical properties. The atomic number of an element is the number of protons in the nucleus of an atom of that element.

21. Atomic Numbers Therefore, the atomic number of an element also tells the number of electrons in a neutral atom of that element. Atoms have no overall electrical charge, an atom must have as many electrons as there are protons in its nucleus.

22. The sum of the protons and neutrons in the nucleus is the mass number of that particular atom. Masses The mass of a neutron is almost the same as the mass of a proton.

23. Masses Isotopes of an element have different mass numbers because they have different numbers of neutrons, but they all have the same atomic number.

24. Atomic Mass In order to have a simpler way of comparing the masses of individual atoms, chemists have devised a different unit of mass called an atomic mass unit, which is given the symbol u. An atom of the carbon-12 isotope contains six protons and six neutrons and has a mass number of 12.

25. Atomic Mass Chemists have defined the carbon-12 atom as having a mass of 12 atomic mass units. Therefore, 1 u = 1/12 the mass of a carbon-12 atom. 1 u is approximately the mass of a single proton or neutron.

26. Calculating Atomic Mass

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

28. Calculating Atomic Mass The atomic mass of Cu-63 is 62.930 amu, and the atomic mass of Cu-65 is 64.928 amu. Use the data above to compute the atomic mass of copper.

29. Calculating Atomic Mass First, calculate the contribution of each isotope to the average atomic mass, being sure to convert each percent to a fractional abundance.

30. Calculating Atomic Mass The average atomic mass of the element is the sum of the mass contributions of each isotope.

31. Question 1 How does the atomic number of an element differ from the element’s mass number? Answer The atomic number of an element is the number of protons in the nucleus. The mass number is the sum of the number of protons and neutrons.

32. Question 2 The table on the next slide shows the five isotopes of germanium found in nature, the abundance of each isotope, and the atomic mass of each isotope. Assessment Questions

33. Calculate the atomic mass of germanium. Assessment Questions Answer 72.59 amu Question 3 IsotopeAbundanceAtomic Mass (amu) Germanium-7021.23% or 0.212369.924 Germanium-7227.66% or 0.276671.922 Germanium-737.73% or 0.077372.923 Germanium-7435.94% or 0.359473.921 Germanium-767.44% or 0.074475.921

34. Information in the Periodic Table The number at the bottom of each box is the average atomic mass of that element. This number is the weighted average mass of all the naturally occurring isotopes of that element.

35. Electrons in Motion Niels Bohr (1885-1962), a Danish scientist who worked with Rutherford, proposed that electrons must have enough energy to keep them in constant motion around the nucleus. Electrons have energy of motion that enables them to overcome the attraction of the positive nucleus.

36. Electrons in Motion Bohr’s view of the atom, which he proposed in 1913, was called the planetary model. This energy keeps the electrons moving around the nucleus.

37. The Electromagnetic Spectrum To boost a satellite into a higher orbit requires energy from a rocket motor. One way to increase the energy of an electron is to supply energy in the form of high- voltage electricity. Another way is to supply electromagnetic radiation, also called radiant energy.

38. The Electron Cloud Model Instead, they are spherical regions of space around the nucleus in which electrons are most likely to be found. As a result of continuing research throughout the 20th century, scientists today realize that energy levels are not neat, planetlike orbits around the nucleus of an atom.

39. The space around the nucleus of an atom where the atom’s electrons are found is called the electron cloud. The Electron Cloud Model These spherical regions where electrons travel may be depicted as clouds around the nucleus. Electrons themselves take up little space but travel rapidly through the space surrounding the nucleus.

40. The Electron Cloud Model

41. Electrons in Energy Level Each energy level can hold a limited number of electrons. How are electrons arranged in energy levels? The lowest energy level is the smallest and the closest to the nucleus.

42. Electrons in Energy Level The second energy level is larger because it is farther away from the nucleus. It holds a maximum of eight electrons. This first energy level holds a maximum of two electrons. The third energy level is larger still and holds a maximum of 18 electrons.

43. Energy Levels A hydrogen atom has only one electron. It’s in the first energy level.

44. Electrons in Energy Level You can also use the periodic table as a tool to predict the number of valence electrons in any atom in Groups 1, 2, 13, 14, 15, 16, 17, and 18. The electrons in the outermost energy level are called valence electrons. All atoms in Group 1, like hydrogen, have one valence electron. Likewise, atoms in Group 2 have two valence electrons.

45. Electrons in Energy Level An oxygen atom has eight electrons. Two of these fill the first energy level, and the remaining six are in the second energy level.

46. Lewis Dot Diagrams Valence electrons are so important to the behavior of an atom, it is useful to represent them with symbols.

47. Lewis Dot Diagrams A Lewis dot diagram illustrates valence electrons as dots (or other small symbols) around the chemical symbol of an element.

48. Lewis Dot Diagrams Each dot represents one valence electron. In the dot diagram, the element’s symbol represents the core of the atom—the nucleus plus all the inner electrons.

49. A. Chlorine Write a Lewis dot diagram for each of the following. Question 2 C. Potassium B. Calcium

50. A. Chlorine Answer C. Potassium B. Calcium

51. A. Ultraviolet light Give an example for each type of electromagnetic energy listed below. Question 3 C. Visible light B. Infrared light

52. Energy Levels and Sublevels Each energy level has a specific number of sublevels, which is the same as the number of the energy level. For example, the first energy level has one sublevel. It’s called the 1s sublevel. The second energy level has two sublevels, the 2s and 2p sublevels

53. Energy Levels and Sublevels Within a given energy level, the energies of the sublevels, from lowest to highest, are s, p, d, and f. The third energy level has three sublevels: the 3s, 3p, and 3d sublevels; and the fourth energy level has four sublevels: the 4s, 4p, 4d, and 4f sublevels.

54. The Distribution of Electrons in Energy Levels A specific number of electrons can go into each sublevel.

55. The Distribution of Electrons in Energy Levels An s sublevel can have a maximum of two electrons, a p sublevel can have six electrons,

56. The Distribution of Electrons in Energy Levels A d sublevel can have ten electrons, and an f sublevel can have 14 electrons.

57. Orbitals This principle is known as the Heisenberg uncertainty principle. In 1932, Heisenberg was awarded the Nobel Prize in Physics for this discovery, which led to the development of the electron cloud model to describe electrons in atoms. In the 1920s, Werner Heisenberg reached the conclusion that it’s impossible to measure accurately both the position and energy of an electron at the same time.

58. Orbitals The electron cloud model is based on the probability of finding an electron in a certain region of space at any given instant. In any atom, electrons are distributed into sublevels and orbitals in the way that creates the most stable arrangement; that is, the one with lowest energy.

59. Electron Configurations This most stable arrangement of electrons in sublevels and orbitals is called an electron configuration. Electrons fill orbitals and sublevels in an orderly fashion beginning with the innermost sublevels and continuing to the outermost.

60. Orbitals and the Periodic Table The periodic table is divided into blocks that show the sublevels and orbitals occupied by the electrons of the atoms. The shape of the modern periodic table is a direct result of the order in which electrons fill energy sublevels and orbitals.

61. Orbitals and the Periodic Table Notice that Groups 1 and 2 (the active metals) have valence electrons in s orbitals, and Groups 13 to 18 (metals, metalloids, and nonmetals) have valence electrons in both s and p orbitals.

62. Building Electron Configurations Chemical properties repeat when elements are arranged by atomic number because electron configurations repeat in a certain pattern. As you move through the table, you’ll notice how an element’s position is related to its electron configuration.

63. Building Electron Configurations This is standard notation for electron configurations. Hydrogen has a single electron in the first energy level. Its electron configuration is 1s 1. The number 1 refers to the energy level, the letter s refers to the sublevel, and the superscript refers to the number of electrons in the sublevel.

64. Building Electron Configurations Helium has two electrons in the 1s orbital. Its electron configuration is 1s 2. Helium has a completely filled first energy level. When the first energy level is filled, additional electrons must go into the second energy level.

65. Building Electron Configurations Lithium begins the second period. Its first two electrons fill the first energy level, so the third electron occupies the second level. Lithium’s electron configuration is 1s 2 2s 1. Beryllium has two electrons in the 2s orbital, so its electron configuration is 1s 2 2s 2.

66. Building Electron Configurations As you continue to move across the second period, electrons begin to enter the p orbitals. Each successive element has one more electron in the 2p orbitals. Carbon, for example, has four electrons in the second energy level. Two of these are in the 2s orbital and two are in the 2p orbitals. The electron configuration for carbon is 1s 2 2s 2 2p 2.

67. Building Electron Configurations

68. At element number 10, neon, the p sublevel is filled with six electrons. The electron configuration for neon is 1s 2 2s 2 2p 6. Neon has eight valence electrons; two are in an s orbital and six are in p orbitals.

69. Building Electron Configurations Notice that neon’s configuration has an inner core of electrons that is identical to the electron configuration in helium (1s 2 ). This insight simplifies the way electron configurations are written. Neon’s electron configuration can be abbreviated: [He]2s 2 2p 6.

70. Building Electron Configurations Notice that elements in the same group have similar configurations. This is important because it shows that the periodic trends in properties, observed in the periodic table, are really the result of repeating patterns of electron configuration.

71. Building Electron Configurations

72. Noble Gases Each period ends with a noble gas, so all the noble gases have filled energy levels and, therefore, stable electron configurations.

73. Noble Gases These stable electron configurations explain the lack of reactivity of the noble gases. Noble gases don’t need to form chemical bonds to acquire stability.

74. Transition Elements Notice in the periodic table that calcium is followed by a group of ten elements beginning with scandium and ending with zinc. These are transition elements. Now the 3d sublevel begins to fill, producing atoms with the lowest possible energy.

75. Transition Elements Like most metals, the transition elements lose electrons to attain a more stable configuration.

76. Inner Transition Elements The two rows beneath the main body of the periodic table are the lanthanides (atomic numbers 58 to 71) and the actinides (atomic numbers 90 to 103). These two series are called inner transition elements because their last electron occupies inner-level 4f orbitals in the sixth period and the 5f orbitals in the seventh period.

77. Assessment Questions Write electron configurations and abbreviated electron configurations of the following elements. Question 1 A. Boron C. Phosphorus B. Fluorine

78. Answer A. Boron C. Phosphorus B. Fluorine Assessment Questions

79. The Electromagnetic Spectrum Radiant energy travels in the form of waves that have both electrical and magnetic properties. These electromagnetic waves can travel through empty space, as you know from the fact that radiant energy from the sun travels to Earth every day.

80. The Electromagnetic Spectrum As you may already have guessed, electromagnetic waves travel through space at the speed of light, which is approximately 300 million meters per second.

81. The Electromagnetic Spectrum Electromagnetic radiation includes radio waves that carry broadcasts to your radio and TV, microwave radiation used to heat food in a microwave oven, radiant heat used to toast bread, and the most familiar form, visible light. All of these forms of radiant energy are parts of a whole range of electromagnetic radiation called the electromagnetic spectrum.

82. The Electromagnetic Spectrum

83. Electrons and Light The spectrum of light released from excited atoms of an element is called the emission spectrum of that element.

84. Evidence for Energy Levels Bohr theorized that electrons absorbed energy and moved to higher energy states. Then, these excited electrons gave off that energy as light waves when they fell back to a lower energy state.

85. Evidence for Energy Levels These regions of space in which electrons can move about the nucleus of an atom are called energy levels. Electrons can have only certain amounts of energy, Bohr reasoned, they can move around the nucleus only at distances that correspond to those amounts of energy.

86. Energy Levels and Sublevels The emission spectrum for each element has a characteristic set of spectral lines. This means that the energy levels within the atom must also be characteristic of each element. But when scientists investigated multi- electron atoms, they found that their spectra were far more complex than would be anticipated by the simple set of energy levels predicted for hydrogen.

87. Energy Levels and Sublevels Notice that these spectra have many more lines than the spectrum of hydrogen.

88. Energy Levels and Sublevels Some lines are grouped close together, and there are big gaps between these groups of lines.

89. Energy Levels and Sublevels The big gaps correspond to the energy released when an electron jumps from one energy level to another.

90. This suggests that sublevels—divisions within a level—exist within a given energy level. Energy Levels and Sublevels The interpretation of the closely spaced lines is that they represent the movement of electrons from levels that are not very different in energy.

91. Sample answers: A. ultraviolet light: Answer the spectrum of light we see as color C. visible light: B. infrared light: part of sunlight radiant heat

92. Energy Levels and Sublevels If electrons are distributed over one or more sublevels within an energy level, then these electrons would have only slightly different energies. The energy sublevels are designated as s, p, d, or f.