Emission and Absorption Spectra
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AP Physics 2 › Emission and Absorption Spectra
A student observes an emission spectrum with discrete lines from a hot, low-density gas. Which statement best explains why the spectrum is not continuous?
The gas emits continuously, but the spectroscope blocks most wavelengths, leaving only a few visible lines.
Electrons emit photons when they absorb energy, producing a few discrete wavelengths instead of a continuum.
Electrons occupy quantized energy levels, so photon energies correspond to specific transitions between levels.
Electrons can have any energy, but they radiate only at certain preferred wavelengths set by intensity.
Explanation
This question tests understanding of emission and absorption spectra. Hot gas atoms have electrons in various excited states due to thermal collisions. These electrons occupy specific quantized energy levels, and when they drop to lower levels, they emit photons with energies equal to the differences between those levels. Since only certain energy differences are allowed, only discrete wavelengths appear, creating a line spectrum rather than a continuum. Choice D incorrectly claims electrons emit photons when absorbing energy, confusing the emission process with absorption. The fundamental principle is: quantized energy levels produce discrete spectral lines, while continuous energy distributions produce continuous spectra.
A low-pressure hydrogen tube emits light that, through a spectroscope, shows several narrow colored lines rather than a rainbow. Which statement best explains the discrete lines?
The electrons emit energy when they absorb photons, producing only a few allowed colors.
The atoms emit photons only when electrons drop between specific energy levels, so only certain photon energies occur.
The spacing of the lines is determined mainly by the intensity of the discharge current in the tube.
The atoms have continuous energy levels, so the spectrum should appear as separated bands instead of a rainbow.
Explanation
This question tests understanding of emission and absorption spectra. In hydrogen atoms, electrons occupy specific quantized energy levels, not continuous ones. When electrons drop from higher to lower energy levels, they emit photons with energies exactly equal to the difference between those levels. Since only certain energy differences are possible, only specific photon energies (and thus specific colors) appear as discrete lines. Choice B incorrectly claims continuous energy levels, which would produce a continuous spectrum, not discrete lines. The key strategy is: discrete spectral lines always indicate quantized energy levels with specific allowed transitions.
Two different gases produce different sets of emission lines under identical conditions. Which statement best explains why the line patterns differ?
Different gases have the same energy levels, but their line spacing changes because their brightness differs.
Different gases have different quantized energy-level spacings, so their allowed transition photon energies differ.
Different gases have continuous energy levels, so random collisions select different wavelengths each time.
Different gases emit photons when absorbing them, so the observed lines depend mainly on the input light intensity.
Explanation
This question tests understanding of emission and absorption spectra. Each element has a unique set of quantized energy levels determined by its nuclear charge and electron configuration. Different gases therefore have different energy-level spacings, leading to different sets of allowed transitions and thus different emission line patterns. This uniqueness allows spectroscopy to identify elements. Choice D incorrectly claims all gases have the same energy levels, which would make spectroscopic identification impossible. The key concept is: each element's unique energy-level structure produces a characteristic spectral fingerprint.
In an emission spectrum, one line corresponds to a photon energy of $3.0\ \text{eV}$. Which statement best describes the atomic process producing this line?
The atom emits a $3.0\ \text{eV}$ photon because its energy levels form a continuum with that energy available.
An electron rises to a higher allowed energy level, releasing a photon of $3.0\ \text{eV}$.
The atom emits a $3.0\ \text{eV}$ photon because the line intensity sets the energy spacing between levels.
An electron drops from a higher to a lower allowed energy level, releasing a photon of $3.0\ \text{eV}$.
Explanation
This question tests understanding of emission and absorption spectra. An emission line at 3.0 eV indicates that an electron has dropped from a higher to a lower quantized energy level, with the energy difference being exactly 3.0 eV. During this downward transition, the electron releases a photon carrying away this energy difference. This is the fundamental process behind all emission spectra. Choice B incorrectly states that electrons rise to higher levels while emitting photons, which violates energy conservation since emission releases energy. The strategy is: emission always involves electron transitions from higher to lower energy levels, releasing photons.
White light passes through cool sodium vapor and the transmitted spectrum contains two dark lines. Which statement best explains the absorption lines?
Atoms absorb photons of all energies, but re-emit only two colors, leaving two dark lines.
Atoms absorb only photons that match energy-level gaps, exciting electrons to higher allowed energy levels.
Electrons emit photons while absorbing them, so the missing colors correspond to emitted energies.
Energy levels are continuous, but the detector resolves only two wavelengths because the light is dim.
Explanation
This question tests understanding of emission and absorption spectra. When white light passes through cool sodium vapor, sodium atoms absorb only those photons whose energies exactly match the energy differences between their quantized levels. These absorbed photons excite electrons from lower to higher energy levels, removing those specific wavelengths from the transmitted light and creating dark absorption lines. The two dark lines correspond to two specific allowed transitions in sodium atoms. Choice B incorrectly suggests atoms absorb all energies but re-emit only two, which violates energy conservation and quantization. The strategy is: absorption lines occur at the same wavelengths as emission lines because they involve the same energy-level transitions.
A student claims an emission spectrum is discrete because atoms emit only certain intensities of light. Which statement best corrects the claim?
The spectrum is discrete because atomic energy levels are continuous, but only certain intensities can be detected.
The spectrum is discrete because higher intensity light creates larger gaps between energy levels in the atom.
The spectrum is discrete because electrons transition between quantized energy levels, fixing photon energies, not intensities.
The spectrum is discrete because electrons emit energy while absorbing photons, limiting the possible emitted intensities.
Explanation
This question tests understanding of emission and absorption spectra. The student confuses intensity (brightness) with energy (color/wavelength). Emission spectra are discrete because electrons transition between quantized energy levels, producing photons with specific energies and thus specific wavelengths/colors. The intensity of a line indicates how many photons are emitted, not their energy. Choice D incorrectly suggests intensity affects energy-level spacing, mixing up these distinct concepts. The crucial distinction is: spectral lines are discrete in wavelength/energy due to quantization, while intensity varies with the number of atoms emitting.
A cool gas in front of a hot continuum source produces dark absorption lines at specific wavelengths. Which statement best explains why only specific wavelengths are removed?
Photons are removed based on their intensity, so the weakest colors disappear as dark lines.
Only photons with energies matching allowed energy-level differences are absorbed to raise electrons to higher levels.
Energy levels are continuous, but collisions remove photons at random wavelengths that appear as lines.
Electrons emit photons during absorption, so the missing wavelengths correspond to emitted photons leaving the beam.
Explanation
This question tests understanding of emission and absorption spectra. Cool gas atoms in their ground or low-energy states can absorb photons from the background continuum source, but only if the photon energies exactly match allowed transitions to higher energy levels. Photons with other energies pass through unaffected because they cannot cause any allowed transitions. This selective absorption creates dark lines at specific wavelengths in the otherwise continuous spectrum. Choice D incorrectly suggests electrons emit photons during absorption, which would violate energy conservation. The key insight is: absorption is wavelength-selective because only photons matching energy-level differences can be absorbed.
A gas absorbs photons of energy $E$ and later emits photons of the same energy $E$ in random directions. Which statement best links this to energy levels?
A photon of energy $E$ matches a specific energy-level gap, so excitation and de-excitation involve the same gap.
The atom’s energy levels form a continuum, so any absorbed energy $E$ can be emitted at the same $E$.
Electrons emit energy while absorbing photons, so the emitted photon energy equals the absorbed energy by conservation.
The emitted photon energy is set by the brightness of the source, so matching energies occur only at high intensity.
Explanation
This question tests understanding of emission and absorption spectra. When a photon of energy E is absorbed, it must match a specific energy gap between quantized levels, exciting an electron from a lower to higher level. When the electron later drops back down, it releases a photon with the same energy E, corresponding to the same energy gap. This process explains resonance fluorescence where absorbed and emitted photons have identical energies. Choice C incorrectly states that electrons emit energy while absorbing photons, which is physically impossible and contradicts energy conservation. The strategy is: absorption and emission involve the same energy-level transitions, so photon energies must match.
White light passes through cool sodium vapor, and dark lines appear in the transmitted spectrum. The absorption line occurs because the atoms
emit photons while absorbing them, canceling most colors except a few wavelengths.
absorb photons of any energy, but only certain wavelengths are removed by collisions.
absorb more strongly at wavelengths where the light source intensity is lowest.
absorb photons whose energies match allowed electron transitions between quantized energy levels.
Explanation
This question tests understanding of emission and absorption spectra. When white light passes through cool sodium vapor, atoms absorb photons whose energies exactly match the energy differences between their quantized electron energy levels. This absorption removes specific wavelengths from the continuous spectrum, creating dark lines at those wavelengths. The absorbed energy promotes electrons from lower to higher energy levels. Choice B incorrectly suggests atoms can absorb photons of any energy, which would create a continuous absorption rather than discrete lines. Remember that absorption lines occur at wavelengths corresponding to allowed electron transitions between quantized energy levels.
In a cool gas, an electron absorbs a photon and moves to a higher energy level. The absorbed photon’s energy must equal
the energy the electron emits while it is being raised to a higher level.
a value set mainly by the intensity of the incident light rather than level spacing.
the difference between two allowed quantized energy levels of the atom.
any value, because electron energy levels form a continuum in a gas.
Explanation
This question tests understanding of emission and absorption spectra. In atoms, electrons occupy quantized energy levels and can only transition between these discrete states. For an electron to move from a lower to a higher energy level, it must absorb a photon whose energy exactly equals the difference between those two levels. If the photon energy doesn't match any allowed transition, it cannot be absorbed and passes through the gas. Choice C incorrectly suggests electrons emit energy while being raised to higher levels, which violates energy conservation. Remember that absorption requires photon energies to exactly match allowed energy-level differences.