Absorption of Light: Videos & Practice Problems

Absorption occurs when an electron in an atom takes in energy, moving to a higher energy state, while emission is the release of that energy as the electron returns to its ground state. The transitions between energy levels produce distinct emission and absorption spectra, identifiable by their unique color patterns. The Lyman series emits UV light, the Balmer series emits visible light, and the Paschen series emits infrared light. The potential energy of electrons can be calculated using the Rydberg constant, $R^{\mathrm{ydberg}}=-1.810-18J(1/n^2)$.

Atomic Emission

The concepts of emission and absorption are fundamental in understanding atomic behavior and energy transitions. Absorption occurs when an electron or atom takes in excess energy, allowing it to move to a higher energy state, often referred to as an excited state. For instance, in Bohr's model of the hydrogen atom, an electron in the first shell can absorb energy from its surroundings, promoting it to a higher shell. However, this excess energy cannot be retained indefinitely; eventually, the electron will release this energy and return to its ground state.

When discussing emissions, we can refer to emission line spectra, which are produced when an electron transitions from a higher energy shell back to a lower one. The periodic table currently recognizes seven electron shells, but this number may increase as new elements are discovered or synthesized. Each transition corresponds to a specific series of emissions. For example, when an electron drops from a higher shell to the first shell, it produces the Lyman series, which falls within the ultraviolet (UV) spectrum. The Balmer series occurs when an electron transitions to the second shell, emitting visible light. Other series, such as the Paschen and Brackett series, correspond to transitions to the third and fourth shells, respectively, and fall within the infrared (IR) spectrum. The Humphreys series is associated with transitions closer to the microwave region.

The energy associated with these transitions can be quantified using the Rydberg formula for potential energy, expressed as:

\( E = -1.8 \times 10^{-18} \, \text{J} \times \frac{1}{n^2} \)

where \( n \) represents the principal quantum number of the shell. As the shell number increases, the distance between shells decreases, which affects the energy released during transitions. A larger drop from a higher shell to a lower shell results in greater energy release.

Emission and absorption spectra are crucial for identifying elements. An emission spectrum displays bright lines against a dark background, representing the frequencies of electromagnetic radiation emitted as electrons fall to lower energy states. Conversely, an absorption spectrum features dark lines on a colored background, indicating the specific wavelengths absorbed by electrons as they transition to higher energy states. The key difference lies in the source of energy: emission spectra originate from excited atoms, while absorption spectra result from light passing through a medium containing atoms that absorb specific wavelengths.

In summary, absorption involves the intake of energy to elevate an electron to a higher state, while emission is the release of that energy as the electron returns to its ground state. Each transition corresponds to a specific series within the electromagnetic spectrum, highlighting the intricate relationship between energy levels and atomic structure.

Here’s what students ask on this topic:

What is the difference between absorption and emission spectra?

Absorption and emission spectra are two ways to analyze the interaction of light with matter. In an absorption spectrum, light passes through a sample, and certain wavelengths are absorbed by the atoms or molecules, resulting in dark lines on a colored background. This indicates the wavelengths of light absorbed as electrons move to higher energy states. In contrast, an emission spectrum is produced when excited electrons return to lower energy states, emitting light at specific wavelengths. This results in bright lines on a dark background. Both spectra are unique to each element and can be used to identify substances.

What is the Lyman series in the hydrogen atom?

The Lyman series refers to the set of spectral lines in the ultraviolet region that are produced when an electron in a hydrogen atom falls from a higher energy level (n > 1) to the ground state (n = 1). These transitions release energy in the form of UV light. The Lyman series is significant in spectroscopy as it helps in understanding the energy levels of hydrogen and the behavior of electrons in atoms.

How is the potential energy of an electron in an atom calculated?

The potential energy of an electron in an atom can be calculated using the Rydberg formula: $E=R\left(\frac{1}{n^2}\right)$, where $R$ is the Rydberg constant, approximately $-1.8\times 10-18J$, and $n$ is the principal quantum number of the electron's energy level. This formula helps determine the energy associated with an electron at a specific energy level in an atom.

What is the Balmer series and its significance?

The Balmer series consists of spectral lines in the visible light region that are produced when an electron in a hydrogen atom falls from a higher energy level (n > 2) to the second energy level (n = 2). These transitions emit light that is visible to the human eye. The Balmer series is significant because it was the first series of lines discovered in the hydrogen spectrum and provided crucial evidence for the quantized nature of energy levels in atoms, leading to the development of quantum mechanics.

How do emission and absorption spectra help in identifying elements?

Emission and absorption spectra are unique to each element, acting like a fingerprint. In an emission spectrum, the bright lines correspond to specific wavelengths of light emitted by electrons as they fall to lower energy levels. In an absorption spectrum, the dark lines represent wavelengths absorbed by electrons as they move to higher energy levels. By comparing these spectra to known reference spectra, scientists can identify the elements present in a sample. This technique is widely used in fields like astronomy, chemistry, and materials science.