Why is the section of the periodic table labeled as

Problem 1CQ A neon discharge emits a bright reddish-orange spectrum. But a glass tube filled with neon is completely transparent. Why doesn’t the neon in the tube absorb orange and red wavelengths?

Problem 1P Figure 29.2 and Table 29.1 showed the wavelengths of the first four lines in the visible spectrum of hydrogen. a. Determine the Balmer formula n and m values for these wavelengths. b. Predict the wavelength of the fifth line in the spectrum.

Problem 2CQ The two spectra shown in Figure Q29.2 belong to the same element, a fictional Element X. Explain why they are different.

Problem 3CQ Is a spectral line with wavelength 656 nm seen in the absorption spectrum of hydrogen atoms? Why or why not?

Problem 2P The wavelengths in the hydrogen spectrum with m = 1 form a series of spectral lines called the Lyman series. Calculate the wavelengths of the first four members of the series.

Problem 3P The Paschen series is analogous to the Balmer series, but with m = 3. Calculate the wavelengths of the first three members in the Paschen series. What part(s) of the electromagnetic spectrum are these in? Section 29.2 Atoms

Problem 4CQ J. J. Thomson studied the ionization of atoms in collisions with electrons. He accelerated electrons through a potential difference, shot them into a gas of atoms, then used a mass spectrometer to detect any ions produced in the collisions. By using different gases, he found that he could produce singly ionized atoms of all the elements that he tried. When he used higher accelerating voltages, he was able to produce doubly ionized atoms of all elements except hydrogen. a. Why did Thomson have to use higher accelerating voltages to detect doubly ionized atoms than to detect singly ionized atoms? b. What conclusion or conclusions about hydrogen atoms can you draw from these observations? Be specific as to how your conclusions are related to the observations.

Problem 4P How many electrons, protons, and neutrons are contained in the following atoms or ions: ?

Problem 5CQ Bohr did not include the gravitational force in his analysis of the hydrogen atom. Is this one of the reasons that his model of the hydrogen atom had only limited success? Explain.

Problem 5P How many electrons, protons, and neutrons are contained in the following atoms or ions: ?

Problem 6CQ If an electron is in a stationary state of an atom, is the electron at rest? If not, what does the term mean?

Problem 6P Write the symbol for an atom or ion with: a. four electrons, four protons, and five neutrons. b. six electrons, seven protons, and eight neutrons.

Problem 7CQ The n = 3 state of hydrogen has . a. Why is the energy negative? b. What is the physical significance of the specific number 1.51 eV?

Problem 7P Write the symbol for an atom or ion with: a. three electrons, three protons, and five neutrons. b. five electrons, six protons, and eight neutrons.

Problem 8CQ For a hydrogen atom, list all possible states that have E = -1.51 eV.

Problem 9CQ What are the n and l values of the following states of a hydrogen atom: (a) 4d, (b) 5f, (c) 6s?

Problem 8P Figure P29.8 is an energy level diagram for a simple atom. What wavelengths appear in the atom’s (a) emission spectrum and (b) absorption spectrum?

Problem 10CQ How would you label the hydrogen-atom states with the following (n, l, m) quantum numbers: (a) (4, 3, 0), (b) (3, 2, 1), (c) (3, 2, -1)?

Problem 9P An electron with 2.0 eV of kinetic energy collides with the atom whose energy-level diagram is shown in Figure P29.8. a. Is the electron able to kick the atom to an excited state? Why or why not? b. If your answer to part a was yes, what is the electron’s kinetic energy after the collision?

Problem 10P The allowed energies of a simple atom are 0.0 eV, 4.0 eV, and 6.0 eV. a. Draw the atom’s energy-level diagram. Label each level with the energy and the principal quantum number. b. What wavelengths appear in the atom’s emission spectrum? c. What wavelengths appear in the atom’s absorption spectrum?

Problem 11CQ Consider the two hydrogen-atom states 5d and 4f. Which has the higher energy? Explain.

Problem 11P The allowed energies of a simple atom are 0.0 eV, 4.0 eV, and 6.0 eV. An electron traveling at a speed of collisionally excites the atom. What are the minimum and maximum speeds the electron could have after the collision?

Problem 12CQ Does each diagram in Figure 12 represent a possible electron configuration of a neutral element? If so, (i) identify the element and (ii) determine if this is the ground state or an excited state. If not, why not? (a) (b) (c) FIGURE 12

Problem 12P A researcher observes hydrogen emitting photons of energy 1.89 eV. What are the quantum numbers of the two states involved in the transition that emits these photons?

Problem 13CQ Do the following electron configurations represent a possible state of an element? If so, (i) identify the element and (ii) determine if this is the ground state or an excited state. If not, why not?

Problem 13P A hydrogen atom is in the n = 3 state. In the Bohr model, how many electron wavelengths fit around this orbit?

Problem 14CQ Why is the section of the periodic table labeled as “transition elements” exactly 10 elements wide in all rows?

Problem 14P A hydrogen atom is in its n = 1 state. In the Bohr model, what is the ratio of its kinetic energy to its potential energy?

Problem 15CQ An electron is in an f state. Can it undergo a quantum jump to an s state? A p state? A d state? Explain.

Problem 15P Show, by actual calculation, mat me Bohr radius is 0.0529 nm and that the ground-state energy of hydrogen is ?13.60 eV.

Problem 16CQ Figure Q29.17 shows the energy-level diagram of Element X. a. What is the ionization energy of Element X? b. An atom in the ground state absorbs a photon, then emits a photon with a wavelength of 1240 nm. What conclusion can you draw about the energy of the photon that was absorbed? c. An atom in the ground state has a collision with an electron, then emits a photon with a wavelength of 1240 nm. What conclusion can you draw about the initial kinetic energy of the electron?

Problem 16P a. What quantum number of the hydrogen atom comes closest to giving a 500-nm-diameter electron orbit? b. What are the electron’s speed and energy in this state?

Problem 17CQ a. Which states of a hydrogen atom can be excited by a collision with an electron with kinetic energy K = 12.5 eV? Explain. b. After the collision the atom is not in its ground state. What happens to the electron? (i) It bounces off with K > 12.5 eV, (ii) It bounces off with K = 12.5 eV, (iii) It bounces off with K < 12.5 eV, (iv) It is absorbed by the atom. Explain your choice. c. After the collision, the atom emits a photon. List all the possible n? m transitions that might occur as a result of this collision.

Problem 17P a. Calculate the de Broglie wavelength of the electron in the n = 1, 2, and 3 states of the hydrogen atom. Use the information in Table 29.2. b. Show numerically that the circumference of the orbit for each of these stationary states is exactly equal to n de Broglie wavelengths. c. Sketch the de Broglie standing wave for the n = 3 orbit.

Problem 18CQ What is an atom’s ionization energy? In other words, if you know the ionization energy of an atom, what is it that you know about the atom?

Problem 18P Show, by calculation, that the first three states of the hydrogen atom have angular momenta %2?,and 3?,respectively.

Problem 19CQ Figure Q29.20 shows the energy levels of a hypothetical atom. a. What minimum kinetic energy (in eV) must an electron have to collisionally excite this atom and cause the emission of a 620 nm photon? Explain. b. Can an electron with K = 6 eV cause the emission of 620 nm light? If so, what is the final kinetic energy of the electron? If not, why not? c. Can a 6 eV photon cause the emission of 620 nm light from this atom? Why or why not? d. Can a 7 eV photon cause the emission of 620 nm light from this atom? Why or why not?

Problem 19P Determine all possible wavelengths of photons that can be emitted from the n = 4 state of a hydrogen atom.

Problem 20CQ Seven possible transitions are identified on the energy-level diagram in Figure Q29.21. For each, is this an allowed transition? If allowed, is it an emission or an absorption transition, and is the photon infrared, visible, or ultraviolet? If not allowed, why not?

Problem 20P List the quantum numbers of (a) all possible 3p states and (b) all possible 3d states.

Problem 21CQ A 2.0 eV photon is incident on an atom in the p state, as shown in the energy-level diagram in Figure Q29.22. Does the atom undergo an absorption transition, a stimulated emission transition, or neither? Explain.

Problem 21P When all quantum numbers are considered, how many different quantum states are there for a hydrogen atom with n = 1? With n = 2? With n = 3? List the quantum numbers of each state.

Problem 22CQ A glass tube contains , some of which are in the ground state and some of which are excited. Figure Q29.23 shows the populations for the atoms’ three energy levels. Is it possible for these atoms to be a laser? If so, on which transition would laser action occur? If not, why not?

Problem 22P What is the angular momentum of a hydrogen atom in (a) a 4p state and (b) a 5f state? Give your answers as a multiple of .

Problem 23P A hydrogen atom has orbital angular momentum 3.65 × 10?34J · s. a. What letter (s, p, d,or f) describes the electron? b. What is the atom’s minimum possible energy? Explain.

Problem 23MCQ An electron collides with an atom in its ground state. The atom then emits a photon of energy . In this process the change in the electron’s energy is

Problem 24MCQ How many states are in the l = 4 subshell? A. 8 B. 9 C. 16 D. 18 E. 22

Problem 24P A hydrogen atom is in the 5p state. Determine (a) its energy, (b) its angular momentum, (c) its quantum number l, and (d) the possible values of its magnetic quantum number m.

Problem 25MCQ What is the ground-state electron configuration of calcium (Z = 20)?

Problem 25P The angular momentum of a hydrogen atom is . What is the minimum energy, in eV, that this atom could have?

Problem 26MCQ An atom emits a photon with a wavelength of 275 nm. By how much does the atom’s energy change? A. 0.72 Ev B. 1.06 eV C. 2.29 eV D. 3.06 eV E. 4.51 eV

Problem 26P Predict the ground-state electron configurations of Mg, Sr, and Ba.

Problem 27MCQ The energy of a hydrogen atom is -3.40 eV. What is the electron’s kinetic energy? A. 1.70 eV B. 2.62 eV C. 3.40 eV D. 5.73 eV E. 6.80 eV

Problem 27P Predict the ground-state electron configurations of Si, Ge, and Pb.

Problem 28MCQ The angular momentum of an electron in a Bohr hydrogen atom is . What is the atom’s energy? A. -13.60 eV B. -6.73 eV C. -3.40 eV D. -1.51 eV E. -0.47 eV

Problem 28P Identify the element for each of these electron configurations. Then determine whether this configuration is the ground state or an excited state.

Problem 29MCQ A “soft x-ray” photon with an energy of 41.8 eV is absorbed by a hydrogen atom in its ground state, knocking the atom’s electron out. What is the speed of the electron as it leaves the atom?

Problem 29P a. With what element is the 3s subshell first completely filled? b. With what element is the Ad subshell first half-filled?

Problem 30P Identify the element for each of these electron configurations. Then determine whether this configuration is the ground state or an excited state.

Problem 31P Explain what is wrong with these electron configurations:

Problem 32P Which has higher energy: an electron in the 4f state or an electron in the 7s state? Explain.

Problem 33P An electron with a speed of collides with an atom. The collision excites the atom from its ground state (0 eV) to a state with an energy of 3.80 eV. What is the speed of the electron after the collision?

Problem 34P Hydrogen gas absorbs light of wavelength 103 nm. Afterward, what wavelengths are seen in the emission spectrum?

Problem 35P What is the longest wavelength of light that can excite the 4s state of sodium?

Problem 36P An electron with a kinetic energy of 3.90 eV collides with a sodium atom. What possible wavelengths of light are subsequently emitted?

Problem 37P a. Is a 4p? 4s transition allowed in sodium? If so, what is its wavelength? If not, why not? b. Is a 3d? 4s transition allowed in sodium? If so, what is its wavelength? If not, why not?

Problem 38P Figure P29.38 shows a molecular energy-level diagram.What are the longest and shortest wavelengths in) the molecule?s absorption spectrum and (b) the molecule's fluorescence spectrum?

Problem 39P The molecule whose energy-level diagram is shown in Figure P29.38 is illuminated by 2.7 eV photons. What is the longest wavelength of light that the molecule can emit?

Problem 40P A 1000 W carbon dioxide laser emits an infrared laser beam with a wavelength of 10.6 ?m. How many photons are emitted per second?

Problem 41P A 1.00 mW helium-neon laser emits a visible laser beam with a wavelength of 633 nm. How many photons are emitted per second?

Problem 42P In LASIK surgery, a laser is used to reshape the cornea of the eye to improve vision. The laser produces extremely short pulses of light, each containing 1.0 mJ of energy. a. In each pulse there are . What is the wavelength of the laser? b. Each pulse lasts only 20 ns. What is the average power delivered to the eye during a pulse?

Problem 43P A ruby laser emits an intense pulse of light that lasts a mere 10 ns. The light has a wavelength of 690 nm, and each pulse has an energy of 500 mJ. a. How many photons are emitted in each pulse? b. What is the rate of photon emission, in photons per second, during the 10 ns that the laser is “on”?

Problem 44GP A 2.55 eV photon is emitted from a hydrogen atom. What are the Balmer formula n and m values corresponding to this emission?

Problem 45GP Two of the wavelengths emitted by a hydrogen atom are 102.6 nm and 1876 nm. a. What are the Balmer formula n and m values for each of these wavelengths? b. For each of these wavelengths, is the light infrared, visible, or ultraviolet?

Problem 46GP In Example 29.2 it was assumed that the initially stationary gold nucleus would remain motionless during a head-on collision with an 8.3 MeV alpha particle. What is the actual recoil speed of the gold nucleus after that elastic collision? Assume that the mass of a gold nucleus is exactly 50 times the mass of an alpha particle. Hint: Review the discussion of perfectly elastic collisions in Chapter 10.

Problem 47GP Consider the gold isotope . a. How many electrons, protons, and neutrons are in a neutral atom? b. The gold nucleus has a diameter of 14.0 fm. What is the density of matter in a gold nucleus? c. The density of lead is . How many times the density of lead is your answer to part b?

Problem 48GP Consider the lead isotope . a. How many electrons, protons, and neutrons are in a neutral atom? b. The lead nucleus has a diameter of 14.2 fm. What are the electric potential and the electric field strength at the surface of a lead nucleus?

Problem 49GP The diameter of an atom is and the diameter of its nucleus is . What percent of the atom’s volume is occupied by mass and what percent is empty space?

Problem 50GP The charge-to-mass ratio of a nucleus, in units of e/u, is q/m = Z/A. For example, a hydrogen nucleus has q/m = 1/1 = 1. a. Make a graph of charge-to-mass ratio versus proton number Z for nuclei with Z = 5, 10, 15, 20,c, 90. For A, use the average atomic mass shown on the periodic table of elements in Appendix B. Show each of these 18 nuclei as a dot, but don’t connect the dots together as a curve. b. Describe any trend that you notice in your graph. c. What’s happening in the nuclei that is responsible for this trend?

Problem 51GP If the nucleus is a few fm in diameter, the distance between the centers of two protons must be ?2 fm. a. Calculate the repulsive electric force between two protons that are 2.0 fm apart. b. Calculate the attractive gravitational force between two protons that are 2.0 fm apart. Could gravity be the force that holds the nucleus together?

Problem 52GP In a head-on collision, the closest approach of a 6.24 MeV alpha particle to the center of a nucleus is 6.00 fm. The nucleus is in an atom of what element? Assume that the nucleus is heavy enough to remain stationary during the collision.

Problem 53GP A 20 MeV alpha particle is fired toward a nucleus. It follows the path shown in Figure P29.53. What is the alpha particle’s speed when it is closest to the nucleus, 20 fm from its center? Assume that the nucleus doesn’t move.

Problem 54GP The oxygen nucleus has a radius of 3.0 fm. a. With what speed must a proton be fired toward an oxygen nucleus to have a turning point 1.0 fm from the surface? Assume that the nucleus is heavy enough to remain stationary during the collision. b. What is the proton’s kinetic energy in MeV?

Problem 55GP The absorption spectrum of an atom consists of the wavelengths 200 nm, 300 nm, and 500 nm. a. Draw the atom’s energy-level diagram. b. What wavelengths are seen in the atom’s emission spectrum?

Problem 56GP The first three energy levels of the fictitious element X are shown in Figure P29.56. a. What wavelengths are observed in the absorption spectrum of element X? Give your answers in nm. b. State whether each of your wavelengths in part a corresponds to ultraviolet, visible, or infrared light. c. An electron with a speed of collides with an atom of element X. Shortly afterward, the atom emits a 1240 nm photon. What was the electron’s speed after the collision? Assume that, because the atom is so much more massive than the electron, the recoil of the atom is negligible. Hint: The energy of the photon is not the energy transferred to the atom in the collision.

Problem 57GP A simple atom has four lines in its absorption spectrum. Ignoring any selection rules, how many lines will it have in its emission spectrum?

Problem 58GP A simple atom has only two absorption lines, at 250 nm and 600 nm. What is the wavelength of the one line in the emission spectrum that does not appear in the absorption spectrum?

Problem 59GP What is the wavelength of the series limit (i.e., the shortest possible wavelength) of the Lyman series in hydrogen?

Problem 60GP What is the energy of a Bohr hydrogen atom with a 5.18 nm diameter?

Problem 61GP A hydrogen atom in the ground state absorbs a 12.75 eV photon. Immediately after the absorption, the atom undergoes a quantum jump to the next-lowest energy level. What is the wavelength of the photon emitted in this quantum jump?

Problem 62GP a. Calculate the orbital radius and the speed of an electron in both the n = 99 and the n = 100 states of hydrogen. b. Determine the orbital frequency of the electron in each of these states. c. Calculate the frequency of a photon emitted in a 100? 99 transition. d. Compare the photon frequency of part c to the average of your two orbital frequencies from part b. By what percent do they differ?

Problem 63GP Two hydrogen atoms collide head-on. The collision brings both atoms to a halt. Immediately after the collision, both atoms emit a 121.6 nm photon. What was the speed of each atom just before the collision?

Problem 64GP A beam of electrons is incident on a gas of hydrogen atoms. a. What minimum speed must the electrons have to cause the emission of 656 nm light from the 3? 2 transition of hydrogen? b. Through what potential difference must the electrons be accelerated to have this speed?

Problem 65GP A hydrogen atom in its fourth excited state emits a photon with a wavelength of 1282 nm. What is the atom’s maximum possible orbital angular momentum after the emission? Give your answer as a multiple of .

Problem 66GP A particular emission line in the hydrogen spectrum has a wavelength of 656.5 nm. What are all possible transitions (e.g., 6d? 2s) that could give rise to this emission?

Problem 67GP a. What downward transitions are possible for a sodium atom in the 6s state? (See Figure 29.24.) b. What are the wavelengths of the photons emitted in each of these transitions?

Problem 68GP The 5d? 3p transition in the emission spectrum of sodium (see Figure 29.24) has a wavelength of 499 nm. What is the energy of the 5d state?

Problem 69GP A sodium atom (see Figure 29.24) emits a photon with wavelength 818 nm shortly after being struck by an electron. What minimum speed did the electron have before the collision?

Problem 70GP Figure P29.72 shows a few energy levels of the mercury atom. One valence electron is always in the 6s state; the other electron changes states. a. Make a table showing all the allowed transitions in the emission spectrum. For each transition, indicate the photon wavelength, in nm. b. What minimum speed must an electron have to excite the 492-nm-wavelength blue emission line in the Hg spectrum?

Problem 71GP Figure P29.73 shows the first few energy levels of the lithium atom. Make a table showing all the allowed transitions in the emission spectrum. For each transition, indicate a. The wavelength, in nm. b. Whether the transition is in the infrared, the visible, or the ultraviolet spectral region. c. Whether or not the transition would be observed in the lithium absorption spectrum.

Problem 72GP A laser emits photons per second from an excited state with energy . a. What is the wavelength of this laser? b. What is the power output of this laser?

Problem 73GP Fluorescence microscopy, discussed in Section 29.8, is an important tool in modern cell biology. A variation on this technique depends on a phenomenon known as two-photon excitation. If two photons are absorbed simultaneously (i.e., within about ), their energies can add. A molecule that is normally excited by a 350 nm photon can be excited by two photons each having half as much energy. For this process to be useful, photons must illuminate the sample at the very high rate of at least . This is achieved by focusing a laser beam to a small spot and by concentrating the power of the laser into very short pulses that are fired 108 times each second. Suppose a biologist wants to use two-photon excitation to excite a molecular species that would be excited by 500 nm light in normal one-photon fluorescence microscopy. What minimum intensity must the laser beam have during each pulse?

Problem 74PP Light-Emitting Diodes Light-emitting diodes, known by the acronym LED, produce the familiar green and red indicator lights used in a wide variety of consumer electronics. LEDs are semiconductor devices in which the electrons can exist only in certain energy levels. Much like molecules, the energy levels are packed together close enough to form what appears to be a continuous band of possible energies. Energy supplied to an LED in a circuit excites electrons from a valence band into a conduction band. An electron can emit a photon by undergoing a quantum jump from a state in the conduction band into an empty state in the valence band, as shown in Figure P29.76. The size of the band gap determines the possible energies— and thus the wavelengths—of the emitted photons. Most LEDs emit a narrow range of wavelengths and thus have a distinct color. This makes them well-suited for traffic lights and other applications where a certain color is desired, but it makes them less desirable for general illumination. One way to make a “white” LED is to combine a blue LED with a substance that fluoresces yellow when illuminated with the blue light. The combination of the two colors makes light that appears reasonably white. An LED emits green light. Increasing the size of the band gap could change the color of the emitted light to A. Red B. Orange C. Yellow D. Blue

Problem 75PP Light-Emitting Diodes Light-emitting diodes, known by the acronym LED, produce the familiar green and red indicator lights used in a wide variety of consumer electronics. LEDs are semiconductor devices in which the electrons can exist only in certain energy levels. Much like molecules, the energy levels are packed together close enough to form what appears to be a continuous band of possible energies. Energy supplied to an LED in a circuit excites electrons from a valence band into a conduction band. An electron can emit a photon by undergoing a quantum jump from a state in the conduction band into an empty state in the valence band, as shown in Figure P29.76. The size of the band gap determines the possible energies— and thus the wavelengths—of the emitted photons. Most LEDs emit a narrow range of wavelengths and thus have a distinct color. This makes them well-suited for traffic lights and other applications where a certain color is desired, but it makes them less desirable for general illumination. One way to make a “white” LED is to combine a blue LED with a substance that fluoresces yellow when illuminated with the blue light. The combination of the two colors makes light that appears reasonably white. Suppose the LED band gap is 2.5 eV, which corresponds to a wavelength of 500 nm. Consider the possible electron transitions in Figure P29.76. 500 nm is the A. Maximum wavelength of the LED. B. Average wavelength of the LED. C. Minimum wavelength of the LED.

Problem 76PP Light-Emitting Diodes Light-emitting diodes, known by the acronym LED, produce the familiar green and red indicator lights used in a wide variety of consumer electronics. LEDs are semiconductor devices in which the electrons can exist only in certain energy levels. Much like molecules, the energy levels are packed together close enough to form what appears to be a continuous band of possible energies. Energy supplied to an LED in a circuit excites electrons from a valence band into a conduction band. An electron can emit a photon by undergoing a quantum jump from a state in the conduction band into an empty state in the valence band, as shown in Figure P29.76. The size of the band gap determines the possible energies— and thus the wavelengths—of the emitted photons. Most LEDs emit a narrow range of wavelengths and thus have a distinct color. This makes them well-suited for traffic lights and other applications where a certain color is desired, but it makes them less desirable for general illumination. One way to make a “white” LED is to combine a blue LED with a substance that fluoresces yellow when illuminated with the blue light. The combination of the two colors makes light that appears reasonably white. The same kind of semiconducting material used to make an LED can also be used to convert absorbed light into electrical energy, essentially operating as an LED in reverse. In this case, the absorption of a photon causes an electron transition from a filled state in the valence band to an unfilled state in the conduction band. If , what is the minimum wavelength of electromagnetic radiation that could lead to electric energy output? A. 140 nm B. 890 nm C. 1400 nm D. 8900 nm

Problem 77PP Light-Emitting Diodes Light-emitting diodes, known by the acronym LED, produce the familiar green and red indicator lights used in a wide variety of consumer electronics. LEDs are semiconductor devices in which the electrons can exist only in certain energy levels. Much like molecules, the energy levels are packed together close enough to form what appears to be a continuous band of possible energies. Energy supplied to an LED in a circuit excites electrons from a valence band into a conduction band. An electron can emit a photon by undergoing a quantum jump from a state in the conduction band into an empty state in the valence band, as shown in Figure P29.76. The size of the band gap determines the possible energies— and thus the wavelengths—of the emitted photons. Most LEDs emit a narrow range of wavelengths and thus have a distinct color. This makes them well-suited for traffic lights and other applications where a certain color is desired, but it makes them less desirable for general illumination. One way to make a “white” LED is to combine a blue LED with a substance that fluoresces yellow when illuminated with the blue light. The combination of the two colors makes light that appears reasonably white. The efficiency of a light source is the percentage of its energy input that gets radiated as visible light. If some of the blue light in an LED is used to cause a fluorescent material to glow, A. The overall efficiency of the LED is increased. B. The overall efficiency of the LED does not change. C. The overall efficiency of the LED decreases.