1. How to Use This Presentation <ul><li>To View the presentation as a slideshow with effects </li></ul><ul><li>select “View” on the menu bar and click on “Slide Show.” </li></ul><ul><li>To advance through the presentation, click the right-arrow key or the space bar. </li></ul><ul><li>From the resources slide, click on any resource to see a presentation for that resource. </li></ul><ul><li>From the Chapter menu screen click on any lesson to go directly to that lesson’s presentation. </li></ul><ul><li>You may exit the slide show at any time by pressing the Esc key. </li></ul>
3. Table of Contents <ul><li>Section 1 Quantization of Energy </li></ul><ul><li>Section 2 Models of the Atom </li></ul><ul><li>Section 3 Quantum Mechanics </li></ul>Atomic Physics Chapter 21
4. Objectives <ul><li>Explain how Planck resolved the ultraviolet catastrophe in blackbody radiation. </li></ul><ul><li>Calculate energy of quanta using Planck’s equation. </li></ul><ul><li>Solve problems involving maximum kinetic energy, work function, and threshold frequency in the photoelectric effect. </li></ul>Section 1 Quantization of Energy Chapter 21
5. Blackbody Radiation <ul><li>Physicists study blackbody radiation by observing a hollow object with a small opening, as shown in the diagram. </li></ul>Section 1 Quantization of Energy Light enters this hollow object through the small opening and strikes the interior wall. Some of the energy is absorbed, but some is reflected at a random angle. After many reflections, essentially all of the incoming energy is absorbed by the cavity wall. <ul><li>A blackbody is a perfect radiator and absorber and emits radiation based only on its temperature. </li></ul>Chapter 21
6. Blackbody Radiation, continued <ul><li>The ultraviolet catastrophe is the failed prediction of classical physics that the energy radiated by a blackbody at extremely short wavelengths is extremely large and that the total energy radiated is infinite. </li></ul><ul><li>Max Planck (1858–1947) developed a formula for blackbody radiation that was in complete agreement with experimental data at all wavelengths by assuming that energy comes in discrete units, or is quantized. </li></ul>Chapter 21 Section 1 Quantization of Energy
7. Blackbody Radiation Chapter 21 The graph on the left shows the intensity of blackbody radiation at three different temperatures. Classical theory’s prediction for blackbody radiation (the blue curve) did not correspond to the experimental data (the red data points) at all wavelengths, whereas Planck’s theory (the red curve) did. Section 1 Quantization of Energy
8. Blackbody Radiation and the Ultraviolet Catastrophe Chapter 21 Section 1 Quantization of Energy
9. Quantum Energy <ul><li>Einstein later applied the concept of quantized energy to light. The units of light energy called quanta (now called photons ) are absorbed or given off as a result of electrons “jumping” from one quantum state to another. </li></ul>Chapter 21 E = hf energy of a quantum = Planck’s constant frequency Planck’s constant ( h ) ≈ 6.63 10 –34 J•s <ul><li>The energy of a light quantum, which corresponds to the energy difference between two adjacent levels, is given by the following equation: </li></ul>Section 1 Quantization of Energy
10. Quantum Energy <ul><li>If Planck’s constant is expressed in units of J•s , the equation E = hf gives the energy in joules. </li></ul><ul><li>However, in atomic physics, energy is often expressed in units of the electron volt, eV. </li></ul><ul><li>An electron volt is defined as the energy that an electron or proton gains when it is accelerated through a potential difference of 1 V. </li></ul><ul><li>The relation between the electron volt and the joule is as follows: </li></ul><ul><li>1 eV = 1.60 10 –19 J </li></ul>Chapter 21 Section 1 Quantization of Energy
11. Energy of a Photon Chapter 21 Section 1 Quantization of Energy
12. The Photoelectric Effect <ul><li>The photoelectric effect is the emission of electrons from a material surface that occurs when light of certain frequencies shines on the surface of the material. </li></ul><ul><li>Classical physics cannot explain the photoelectric effect. </li></ul><ul><li>Einstein assumed that an electromagnetic wave can be viewed as a stream of particles called photons. Photon theory accounts for observations of the photoelectric effect. </li></ul>Chapter 21 Section 1 Quantization of Energy
13. The Photoelectric Effect Chapter 21 Section 1 Quantization of Energy
14. The Photoelectric Effect Chapter 21 Section 1 Quantization of Energy
15. The Photoelectric Effect, continued <ul><li>No electrons are emitted if the frequency of the incoming light falls below a certain frequency, called the threshold frequency ( f t ). </li></ul><ul><li>The smallest amount of energy the electron must have to escape the surface of a metal is the work function of the metal. </li></ul><ul><li>The work function is equal to hf t . </li></ul>Chapter 21 Section 1 Quantization of Energy
16. The Photoelectric Effect, continued <ul><li>Because energy must be conserved, the maximum kinetic energy (of photoelectrons ejected from the surface) is the difference between the photon energy and the work function of the metal. </li></ul>maximum kinetic energy of a photoelectron KE max = hf – hf t maximum kinetic energy = (Planck’s constant frequency of incoming photon) – work function Chapter 21 Section 1 Quantization of Energy
17. Compton Shift <ul><li>If light behaves like a particle, then photons should have momentum as well as energy; both quantities should be conserved in elastic collisions. </li></ul><ul><li>The American physicist Arthur Compton directed X rays toward a block of graphite to test this theory. </li></ul><ul><li>He found that the scattered waves had less energy and longer wavelengths than the incoming waves, just as he had predicted. </li></ul><ul><li>This change in wavelength, known as the Compton shift, supports Einstein’s photon theory of light. </li></ul>Chapter 21 Section 1 Quantization of Energy
18. Compton Shift Chapter 21 Section 1 Quantization of Energy
19. Objectives <ul><li>Explain the strengths and weaknesses of Rutherford’s model of the atom. </li></ul><ul><li>Recognize that each element has a unique emission and absorption spectrum. </li></ul><ul><li>Explain atomic spectra using Bohr’s model of the atom. </li></ul><ul><li>Interpret energy-level diagrams. </li></ul>Section 2 Models of the Atom Chapter 21
20. Early Models of the Atom <ul><li>The model of the atom in the days of Newton was that of a tiny, hard, indestructible sphere. </li></ul><ul><li>The discovery of the electron in 1897 prompted J. J. Thomson (1856–1940) to suggest a new model of the atom. </li></ul><ul><li>In Thomson’s model, electrons are embedded in a spherical volume of positive charge like seeds in a watermelon. </li></ul>Section 2 Models of the Atom Chapter 21
21. Early Models of the Atom, continued <ul><li>Ernest Rutherford (1871–1937) later proved that Thomson’s model could not be correct. </li></ul><ul><li>In his experiment, a beam of positively charged alpha particles was projected against a thin metal foil. </li></ul>Section 2 Models of the Atom <ul><li>Most of the alpha particles passed through the foil. Some were deflected through very large angles. </li></ul>Chapter 21
22. Rutherford’s Gold Foil Experiment Section 2 Models of the Atom Chapter 21
23. Early Models of the Atom, continued <ul><li>Rutherford concluded that all of the positive charge in an atom and most of the atom’s mass are found in a region that is small compared to the size of the atom. </li></ul><ul><li>He called this region the the nucleus of the atom. </li></ul><ul><li>Any electrons in the atom were assumed to be in the relatively large volume outside the nucleus. </li></ul>Section 2 Models of the Atom Chapter 21
24. Early Models of the Atom, continued <ul><li>To explain why electrons were not pulled into the nucleus, Rutherford viewed the electrons as moving in orbits about the nucleus. </li></ul>Section 2 Models of the Atom <ul><li>However, accelerated charges should radiate electromagnetic waves, losing energy. This would lead to a rapid collapse of the atom. </li></ul><ul><li>This difficulty led scientists to continue searching for a new model of the atom. </li></ul>Chapter 21
25. Atomic Spectra <ul><li>When the light given off by an atomic gas is passed through a prism, a series of distinct bright lines is seen. Each line corresponds to a different wavelength, or color. </li></ul>Section 2 Models of the Atom <ul><li>A diagram or graph that indicates the wavelengths of radiant energy that a substance emits is called an emission spectrum. </li></ul><ul><li>Every element has a distinct emission spectrum. </li></ul>Chapter 21
26. Atomic Spectra, continued <ul><li>An element can also absorb light at specific wavelengths. </li></ul><ul><li>The spectral lines corresponding to this process form what is known as an absorption spectrum. </li></ul><ul><li>An absorption spectrum can be seen by passing light containing all wavelengths through a vapor of the element being analyzed. </li></ul><ul><li>Each line in the absorption spectrum of a given element coincides with a line in the emission spectrum of that element. </li></ul>Section 2 Models of the Atom Chapter 21
27. Emission and Absorption Spectra of Hydrogen Section 2 Models of the Atom Chapter 21
28. The Bohr Model of the Hydrogen Atom <ul><li>In 1913, the Danish physicist Niels Bohr (1885–1962) proposed a new model of the hydrogen atom that explained atomic spectra. </li></ul><ul><li>In Bohr’s model, only certain orbits are allowed. The electron is never found between these orbits; instead, it is said to “jump” instantly from one orbit to another. </li></ul><ul><li>In Bohr’s model, transitions between stable orbits with different energy levels account for the discrete spectral lines. </li></ul>Section 2 Models of the Atom Chapter 21
29. The Bohr Model, continued <ul><li>When light of a continuous spectrum shines on the atom, only the photons whose energy ( hf ) matches the energy separation between two levels can be absorbed by the atom. </li></ul>Section 2 Models of the Atom <ul><li>When this occurs, an electron jumps from a lower energy state to a higher energy state, which corresponds to an orbit farther from the nucleus. </li></ul><ul><li>This is called an excited state . The absorbed photons account for the dark lines in the absorption spectrum. </li></ul>Chapter 21
30. The Bohr Model, continued <ul><li>Once an electron is in an excited state, there is a certain probability that it will jump back to a lower energy level by emitting a photon. </li></ul><ul><li>This process is called spontaneous emission. </li></ul>Section 2 Models of the Atom <ul><li>The emitted photons are responsible for the bright lines in the emission spectrum. </li></ul><ul><li>In both cases, there is a correlation between the “size” of an electron’s jump and the energy of the photon. </li></ul>Chapter 21
31. The Bohr Model of the Atom Section 2 Models of the Atom Chapter 21
32. Sample Problem <ul><li>Interpreting Energy-Level Diagrams </li></ul><ul><li>An electron in a hydrogen atom drops from energy level E 4 to energy level E 2 . What is the frequency of the emitted photon, and which line in the emission spectrum corresponds to this event? </li></ul>Section 2 Models of the Atom Chapter 21
33. Sample Problem, continued <ul><li>Find the energy of the photon. </li></ul><ul><li>The energy of the photon is equal to the change in the energy of the electron. The electron’s initial energy level was E 4 , and the electron’s final energy level was E 2 . Using the values from the energy-level diagram gives the following: </li></ul><ul><li>E = E initia l – E final = E 4 – E 2 </li></ul><ul><li>E = (–0.850 eV) – (–3.40 eV) = 2.55 eV </li></ul>Section 2 Models of the Atom Chapter 21
34. Sample Problem, continued <ul><li>Tip: Note that the energies for each energy level are negative. The reason is that the energy of an electron in an atom is defined with respect to the amount of work required to remove the electron from the atom. In some energy-level diagrams, the energy of E 1 is defined as zero, and the higher energy levels are positive. </li></ul><ul><li>In either case, the difference between a higher energy level and a lower one is always positive, indicating that the electron loses energy when it drops to a lower level . </li></ul>Section 2 Models of the Atom Chapter 21
35. Sample Problem, continued <ul><li>Tip: Note that electron volts were converted to joules so that the units cancel properly. </li></ul>Section 2 Models of the Atom 2. Use Planck’s equation to find the frequency. Chapter 21
36. Sample Problem, continued Section 2 Models of the Atom <ul><li>Find the corresponding line in the emission spectrum. </li></ul>Examination of the diagram shows that the electron’s jump from energy level E 4 to energy level E 2 corresponds to Line 3 in the emission spectrum. Chapter 21
37. Sample Problem, continued Section 2 Models of the Atom 4. Evaluate your answer. Line 3 is in the visible part of the electromagnetic spectrum and appears to be blue. The frequency f = 6.15 10 14 Hz lies within the range of the visible spectrum and is toward the violet end, so it is reasonable that light of this frequency would be visible blue light. Chapter 21
38. <ul><li>Bohr’s model was not considered to be a complete picture of the structure of the atom. </li></ul><ul><ul><li>Bohr assumed that electrons do not radiate energy when they are in a stable orbit, but his model offered no explanation for this. </li></ul></ul><ul><ul><li>Another problem with Bohr’s model was that it could not explain why electrons always have certain stable orbits </li></ul></ul><ul><li>For these reasons, scientists continued to search for a new model of the atom. </li></ul>The Bohr Model, continued Section 2 Models of the Atom Chapter 21
39. Objectives <ul><li>Recognize the dual nature of light and matter. </li></ul><ul><li>Calculate the de Broglie wavelength of matter waves. </li></ul><ul><li>Distinguish between classical ideas of measurement and Heisenberg’s uncertainty principle. </li></ul><ul><li>Describe the quantum-mechanical picture of the atom, including the electron cloud and probability waves. </li></ul>Section 3 Quantum Mechanics Chapter 21
40. The Dual Nature of Light <ul><li>As seen earlier, there is considerable evidence for the photon theory of light. In this theory, all electromagnetic waves consist of photons, particle-like pulses that have energy and momentum. </li></ul><ul><li>On the other hand, light and other electromagnetic waves exhibit interference and diffraction effects that are considered to be wave behaviors. </li></ul><ul><li>So, which model is correct? </li></ul>Section 3 Quantum Mechanics Chapter 21
41. The Dual Nature of Light, continued <ul><li>Some experiments can be better explained or only explained by the photon concept, whereas others require a wave model. </li></ul><ul><li>Most physicists accept both models and believe that the true nature of light is not describable in terms of a single classical picture. </li></ul><ul><ul><li>At one extreme, the electromagnetic wave description suits the overall interference pattern formed by a large number of photons. </li></ul></ul><ul><ul><li>At the other extreme, the particle description is more suitable for dealing with highly energetic photons of very short wavelengths. </li></ul></ul>Section 3 Quantum Mechanics Chapter 21
42. The Dual Nature of Light Section 3 Quantum Mechanics Chapter 21
43. Matter Waves <ul><li>In 1924, the French physicist Louis de Broglie (1892–1987) extended the wave-particle duality. De Broglie proposed that all forms of matter may have both wave properties and particle properties. </li></ul><ul><li>Three years after de Broglie’s proposal, C. J. Davisson and L. Germer, of the United States, discovered that electrons can be diffracted by a single crystal of nickel. This important discovery provided the first experimental confirmation of de Broglie’s theory. </li></ul>Section 3 Quantum Mechanics Chapter 21
44. Matter Waves, continued <ul><li>The wavelength of a photon is equal to Planck’s constant ( h ) divided by the photon’s momentum ( p ). De Broglie speculated that this relationship might also hold for matter waves, as follows: </li></ul>Section 3 Quantum Mechanics <ul><li>As seen by this equation, the larger the momentum of an object, the smaller its wavelength. </li></ul>Chapter 21
45. Matter Waves, continued <ul><li>In an analogy with photons, de Broglie postulated that the frequency of a matter wave can be found with Planck’s equation, as illustrated below: </li></ul>Section 3 Quantum Mechanics <ul><li>The dual nature of matter suggested by de Broglie is quite apparent in the wavelength and frequency equations, both of which contain particle concepts ( E and mv ) and wave concepts ( and f ). </li></ul>Chapter 21
46. Matter Waves, continued <ul><li>He assumed that an electron orbit would be stable only if it contained an integral (whole) number of electron wavelengths. </li></ul>Section 3 Quantum Mechanics <ul><li>De Broglie saw a connection between his theory of matter waves and the stable orbits in the Bohr model. </li></ul>Chapter 21
47. De Broglie and the Wave-Particle Nature of Electrons Section 3 Quantum Mechanics Chapter 21
48. The Uncertainty Principle <ul><li>In 1927, Werner Heisenberg argued that it is fundamentally impossible to make simultaneous measurements of a particle’s position and momentum with infinite accuracy. </li></ul><ul><li>In fact, the more we learn about a particle’s momentum, the less we know of its position, and the reverse is also true. </li></ul><ul><li>This principle is known as Heisenberg’s uncertainty principle. </li></ul>Section 3 Quantum Mechanics Chapter 21
49. The Uncertainty Principle Section 3 Quantum Mechanics Chapter 21
50. The Electron Cloud, continued <ul><li>Quantum mechanics also predicts that the electron can be found in a spherical region surrounding the nucleus. </li></ul><ul><li>This result is often interpreted by viewing the electron as a cloud surrounding the nucleus. </li></ul><ul><li>Analysis of each of the energy levels of hydrogen reveals that the most probable electron location in each case is in agreement with each of the radii predicted by the Bohr theory. </li></ul>Section 3 Quantum Mechanics Chapter 21
51. The Electron Cloud <ul><li>Because the electron’s location cannot be precisely determined, it is useful to discuss the probability of finding the electron at different locations. </li></ul><ul><li>The diagram shows the probability per unit distance of finding the electron at various distances from the nucleus in the ground state of hydrogen. </li></ul>Section 3 Quantum Mechanics Chapter 21
52. Multiple Choice <ul><li>1. What is another word for “quantum of light”? </li></ul><ul><li>A. blackbody radiation </li></ul><ul><li>B. energy level </li></ul><ul><li>C. frequency </li></ul><ul><li>D. photon </li></ul>Standardized Test Prep Chapter 21
53. Multiple Choice <ul><li>1. What is another word for “quantum of light”? </li></ul><ul><li>A. blackbody radiation </li></ul><ul><li>B. energy level </li></ul><ul><li>C. frequency </li></ul><ul><li>D. photon </li></ul>Standardized Test Prep Chapter 21
54. Multiple Choice, continued <ul><li>2. According to classical physics, when a light illuminates a photosensitive surface, what should determine how long it takes before electrons are ejected from the surface? </li></ul><ul><li>F. frequency </li></ul><ul><li>G. intensity </li></ul><ul><li>H. photon energy </li></ul><ul><li>J. wavelength </li></ul>Standardized Test Prep Chapter 21
55. Multiple Choice, continued <ul><li>2. According to classical physics, when a light illuminates a photosensitive surface, what should determine how long it takes before electrons are ejected from the surface? </li></ul><ul><li>F. frequency </li></ul><ul><li>G. intensity </li></ul><ul><li>H. photon energy </li></ul><ul><li>J. wavelength </li></ul>Standardized Test Prep Chapter 21
56. Multiple Choice, continued <ul><li>3. According to Einstein’s photon theory of light, what does the intensity of light shining on a metal determine? </li></ul><ul><li>A. the number of photons hitting the metal in a given time interval </li></ul><ul><li>B. the energy of photons hitting the metal </li></ul><ul><li>C. whether or not photoelectrons will be emitted </li></ul><ul><li>D. KE max of emitted photoelectrons </li></ul>Standardized Test Prep Chapter 21
57. Multiple Choice, continued <ul><li>3. According to Einstein’s photon theory of light, what does the intensity of light shining on a metal determine? </li></ul><ul><li>A. the number of photons hitting the metal in a given time interval </li></ul><ul><li>B. the energy of photons hitting the metal </li></ul><ul><li>C. whether or not photoelectrons will be emitted </li></ul><ul><li>D. KE max of emitted photoelectrons </li></ul>Standardized Test Prep Chapter 21
58. Multiple Choice, continued <ul><li>4. An X-ray photon is scattered by a stationary electron. How does the frequency of this scattered photon compare to its frequency before being scattered? </li></ul><ul><li>F. The new frequency is higher. </li></ul><ul><li>G. The new frequency is lower. </li></ul><ul><li>H. The frequency stays the same. </li></ul><ul><li>J. The scattered photon has no frequency. </li></ul>Standardized Test Prep Chapter 21
59. Multiple Choice, continued <ul><li>4. An X-ray photon is scattered by a stationary electron. How does the frequency of this scattered photon compare to its frequency before being scattered? </li></ul><ul><li>F. The new frequency is higher. </li></ul><ul><li>G. The new frequency is lower. </li></ul><ul><li>H. The frequency stays the same. </li></ul><ul><li>J. The scattered photon has no frequency. </li></ul>Standardized Test Prep Chapter 21
60. Multiple Choice, continued <ul><li>5. Which of the following summarizes Thomson’s model of the atom? </li></ul><ul><li>A. Atoms are hard, uniform, indestructible spheres. </li></ul><ul><li>B. Electrons are embedded in a sphere of positive charge. </li></ul><ul><li>C. Electrons orbit the nucleus in the same way </li></ul><ul><li>that planets orbit the sun. </li></ul><ul><li>D. Electrons exist only at discrete energy levels. </li></ul>Standardized Test Prep Chapter 21
61. Multiple Choice, continued <ul><li>5. Which of the following summarizes Thomson’s model of the atom? </li></ul><ul><li>A. Atoms are hard, uniform, indestructible spheres. </li></ul><ul><li>B. Electrons are embedded in a sphere of positive charge. </li></ul><ul><li>C. Electrons orbit the nucleus in the same way </li></ul><ul><li>that planets orbit the sun. </li></ul><ul><li>D. Electrons exist only at discrete energy levels. </li></ul>Standardized Test Prep Chapter 21
62. Multiple Choice, continued <ul><li>6. What happens when an electron moves from a higher energy level to a lower energy level in an atom? </li></ul><ul><li>F. Energy is absorbed from a source outside the atom. </li></ul><ul><li>G. The energy contained in the electromagnetic field inside the atom increases. </li></ul><ul><li>H. Energy is released across a continuous range of values. </li></ul><ul><li>J. A photon is emitted with energy equal to the difference in energy between the two levels. </li></ul>Standardized Test Prep Chapter 21
63. Multiple Choice, continued <ul><li>6. What happens when an electron moves from a higher energy level to a lower energy level in an atom? </li></ul><ul><li>F. Energy is absorbed from a source outside the atom. </li></ul><ul><li>G. The energy contained in the electromagnetic field inside the atom increases. </li></ul><ul><li>H. Energy is released across a continuous range of values. </li></ul><ul><li>J. A photon is emitted with energy equal to the difference in energy between the two levels. </li></ul>Standardized Test Prep Chapter 21
64. Multiple Choice, continued <ul><li>Use the energy-level diagram for hydrogen to answer </li></ul><ul><li>questions 7–8. </li></ul>Standardized Test Prep Chapter 21 7. What is the frequency of the photon emitted when an electron jumps from E 5 to E 2 ? A. 2.86 eV B. 6.15 10 14 Hz C. 6.90 10 14 Hz D. 4.31 10 33 Hz
65. Multiple Choice, continued <ul><li>Use the energy-level diagram for hydrogen to answer </li></ul><ul><li>questions 7–8. </li></ul>Standardized Test Prep Chapter 21 7. What is the frequency of the photon emitted when an electron jumps from E 5 to E 2 ? A. 2.86 eV B. 6.15 10 14 Hz C. 6.90 10 14 Hz D. 4.31 10 33 Hz
66. Multiple Choice, continued <ul><li>Use the energy-level diagram for hydrogen to answer </li></ul><ul><li>questions 7–8. </li></ul>Standardized Test Prep Chapter 21 8. What frequency of photon would be absorbed when an electron jumps from E 2 to E 3 ? F. 1.89 eV G. 4.56 10 14 Hz H. 6.89 10 14 Hz J. 2.85 10 33 Hz
67. Multiple Choice, continued <ul><li>Use the energy-level diagram for hydrogen to answer </li></ul><ul><li>questions 7–8. </li></ul>Standardized Test Prep Chapter 21 8. What frequency of photon would be absorbed when an electron jumps from E 2 to E 3 ? F. 1.89 eV G. 4.56 10 14 Hz H. 6.89 10 14 Hz J. 2.85 10 33 Hz
68. Multiple Choice, continued <ul><li>9. What type of spectrum is created by applying a high potential difference to a pure atomic gas? </li></ul><ul><li>A. an emission spectrum </li></ul><ul><li>B. an absorption spectrum </li></ul><ul><li>C. a continuous spectrum </li></ul><ul><li>D. a visible spectrum </li></ul>Standardized Test Prep Chapter 21
69. Multiple Choice, continued <ul><li>9. What type of spectrum is created by applying a high potential difference to a pure atomic gas? </li></ul><ul><li>A. an emission spectrum </li></ul><ul><li>B. an absorption spectrum </li></ul><ul><li>C. a continuous spectrum </li></ul><ul><li>D. a visible spectrum </li></ul>Standardized Test Prep Chapter 21
70. Multiple Choice, continued <ul><li>10. What type of spectrum is used to identify elements in the atmospheres of stars? </li></ul><ul><li>F. an emission spectrum </li></ul><ul><li>G. an absorption spectrum </li></ul><ul><li>H. a continuous spectrum </li></ul><ul><li>J. a visible spectrum </li></ul>Standardized Test Prep Chapter 21
71. Multiple Choice, continued <ul><li>10. What type of spectrum is used to identify elements in the atmospheres of stars? </li></ul><ul><li>F. an emission spectrum </li></ul><ul><li>G. an absorption spectrum </li></ul><ul><li>H. a continuous spectrum </li></ul><ul><li>J. a visible spectrum </li></ul>Standardized Test Prep Chapter 21
72. Multiple Choice, continued <ul><li>11. What is the speed of a proton ( m = 1.67 10 –27 kg) with a de Broglie wavelength of 4.00 10 –14 m? </li></ul><ul><li>A. 1.59 10 –30 m/s </li></ul><ul><li>B. 1.01 10 –7 m/s </li></ul><ul><li>C. 9.93 10 6 m/s </li></ul><ul><li>D. 1.01 10 7 m/s </li></ul>Standardized Test Prep Chapter 21
73. Multiple Choice, continued <ul><li>11. What is the speed of a proton ( m = 1.67 10 –27 kg) with a de Broglie wavelength of 4.00 10 –14 m? </li></ul><ul><li>A. 1.59 10 –30 m/s </li></ul><ul><li>B. 1.01 10 –7 m/s </li></ul><ul><li>C. 9.93 10 6 m/s </li></ul><ul><li>D. 1.01 10 7 m/s </li></ul>Standardized Test Prep Chapter 21
74. Multiple Choice, continued <ul><li>12. What does Heisenberg’s uncertainty principle state? </li></ul><ul><li>F. It is impossible to simultaneously measure a particle’s position and momentum with infinite accuracy. </li></ul><ul><li>G. It is impossible to measure both a particle’s position and its momentum. </li></ul><ul><li>H. The more accurately we know a particle’s position, the more accurately we know the particle’s momentum. </li></ul><ul><li>J. All measurements are uncertain. </li></ul>Standardized Test Prep Chapter 21
75. Multiple Choice, continued <ul><li>12. What does Heisenberg’s uncertainty principle state? </li></ul><ul><li>F. It is impossible to simultaneously measure a particle’s position and momentum with infinite accuracy. </li></ul><ul><li>G. It is impossible to measure both a particle’s position and its momentum. </li></ul><ul><li>H. The more accurately we know a particle’s position, the more accurately we know the particle’s momentum. </li></ul><ul><li>J. All measurements are uncertain. </li></ul>Standardized Test Prep Chapter 21
76. Short Response <ul><li>13. What is the energy of a photon of light with frequency f = 2.80 10 14 Hz? Give your answer in both J and eV. </li></ul>Standardized Test Prep Chapter 21
77. Short Response <ul><li>13. What is the energy of a photon of light with frequency f = 2.80 10 14 Hz? Give your answer in both J and eV. </li></ul><ul><li>Answer: 1.86 10 –19 J, 1.16 eV </li></ul>Standardized Test Prep Chapter 21
78. Short Response, continued <ul><li>14. Light of wavelength 3.0 10 –7 m shines on the metals lithium, iron, and mercury, which have work functions of 2.3 eV, 3.9 eV, and 4.5 eV, respectively. Which of these metals will exhibit the photoelectric effect? For each metal that does exhibit the photoelectric effect, what is the maximum kinetic energy of the photoelectrons? </li></ul>Standardized Test Prep Chapter 21
79. Short Response, continued <ul><li>14. Light of wavelength 3.0 10 –7 m shines on the metals lithium, iron, and mercury, which have work functions of 2.3 eV, 3.9 eV, and 4.5 eV, respectively. Which of these metals will exhibit the photoelectric effect? For each metal that does exhibit the photoelectric effect, what is the maximum kinetic energy of the photoelectrons? </li></ul><ul><li>Answers: lithium and iron </li></ul><ul><li>lithium: 1.8 eV, iron: 0.2 eV </li></ul>Standardized Test Prep Chapter 21
80. Short Response, continued <ul><li>15. Identify the behavior of an electron as primarily like a wave or like a particle in each of the following situations: </li></ul><ul><li>a. traversing a circular orbit in a magnetic field </li></ul><ul><li>b. absorbing a photon and being ejected from the surface of a metal </li></ul><ul><li>c. forming an interference pattern </li></ul>Standardized Test Prep Chapter 21
81. Short Response, continued <ul><li>15. Identify the behavior of an electron as primarily like a wave or like a particle in each of the following situations: </li></ul><ul><li>a. traversing a circular orbit in a magnetic field </li></ul><ul><li>b. absorbing a photon and being ejected from the surface of a metal </li></ul><ul><li>c. forming an interference pattern </li></ul><ul><li>Answers: a. particle; b. particle; c. wave </li></ul>Standardized Test Prep Chapter 21
82. Extended Response <ul><li>16. Describe Bohr’s model of the atom. Identify the assumptions that Bohr made that were a departure from those of classical physics. Explain how Bohr’s model accounts for atomic spectra. </li></ul>Standardized Test Prep Chapter 21
83. Extended Response <ul><li>16. Describe Bohr’s model of the atom. Identify the assumptions that Bohr made that were a departure from those of classical physics. Explain how Bohr’s model accounts for atomic spectra. </li></ul><ul><li>Answer: Answers should describe electrons orbiting a nucleus only in discrete energy levels. The model departs from classical physics in that the electrons are only allowed to have certain energies and they do not lose energy simply by moving in an electromagnetic field. Atomic spectra are the result of photons being emitted or absorbed when electrons jump between energy levels. </li></ul>Standardized Test Prep Chapter 21
84. Extended Response, continued <ul><li>17. Electrons are ejected from a surface with speeds ranging up to 4.6 10 5 m/s when light with a wavelength of 625 nm is used. </li></ul><ul><li>a. What is the work function of this surface? </li></ul><ul><li>b. What is the threshold frequency for this surface? </li></ul><ul><li>Show all your work. </li></ul>Standardized Test Prep Chapter 21
85. Extended Response, continued <ul><li>17. Electrons are ejected from a surface with speeds ranging up to 4.6 10 5 m/s when light with a wavelength of 625 nm is used. </li></ul><ul><li>a. What is the work function of this surface? </li></ul><ul><li>b. What is the threshold frequency for this surface? </li></ul><ul><li>Show all your work. </li></ul><ul><li>Answers: a. 1.39 eV </li></ul><ul><li>b. 3.35 10 14 Hz </li></ul>Standardized Test Prep Chapter 21
86. Extended Response, continued <ul><li>18. The wave nature of electrons makes an electron microscope, which uses electrons rather than light, possible. The resolving power of any microscope is approximately equal to the wavelength used. A resolution of approximately 1.0 10 –11 m would be required in order to “see” an atom. </li></ul><ul><li>a. If electrons were used, what minimum kinetic energy of the electrons (in eV) would be required to obtain this degree of resolution? </li></ul><ul><li>b. If photons were used, what minimum photon energy would be required? </li></ul>Standardized Test Prep Chapter 21
87. Extended Response, continued <ul><li>18. The wave nature of electrons makes an electron microscope, which uses electrons rather than light, possible. The resolving power of any microscope is approximately equal to the wavelength used. A resolution of approximately 1.0 10 –11 m would be required in order to “see” an atom. </li></ul><ul><li>a. If electrons were used, what minimum kinetic energy of the electrons (in eV) would be required to obtain this degree of resolution? </li></ul><ul><li>b. If photons were used, what minimum photon energy would be required? </li></ul><ul><li>Answers: a. 1.5 10 4 eV; b. 1.2 10 5 eV </li></ul>Standardized Test Prep Chapter 21
88. Atomic Spectra Section 2 Models of the Atom Chapter 21
Be the first to comment