Line Spectra Of Hydrogen Atom

Line Spectra Of Hydrogen Atom

Vishal kumarUpdated on 02 Jul 2025, 06:25 PM IST

The line spectra of the hydrogen atom reveal distinct patterns of light emitted when electrons transition between energy levels within the atom. These spectral lines, unique to hydrogen, are crucial for understanding atomic structure and quantum mechanics. Observed as discrete lines in the emission spectrum, they provide evidence for the quantized nature of atomic energy levels. In real life, these spectral lines are foundational in fields such as astronomy, where they help identify the composition of distant stars and galaxies, and in spectroscopy, which aids in analyzing chemical compositions and developing new materials. By studying the hydrogen line spectra, we gain insights into atomic behaviour and advance various scientific and industrial applications. In this article, we will discuss the concept of the Line spectra of a Hydrogen atom and solve examples for better concept clarity.

This Story also Contains

  1. Line Spectra of a Hydrogen Atom
  2. Solved Examples Based on Line Spectra of the Hydrogen Atom
  3. Summary

Line Spectra of a Hydrogen Atom

According to Bohr, when an atom makes a transition from a higher energy level to a lower energy level, it emits a photon with energy equal to the energy difference between the initial and final levels. If Ei, is the initial energy of the atom before such a transition, Ef is its final energy after the transition, then conservation of energy gives the energy of the emitted photon.

$
\begin{aligned}
& \mathrm{h} v=\frac{\mathrm{hc}}{\lambda}=E_i-E_{\mathrm{f}} \\
& \frac{h c}{\lambda}=\frac{-13.6}{n_i^2} \mathrm{eV}-\frac{-13.6}{n_f^2} \mathrm{eV}=13.6 \mathrm{eV}\left(\frac{1}{n_f^2}-\frac{1}{n_i^2}\right) \\
& R c h=13.6 \mathrm{eV}=1 \text { Rydberg energy } \\
& \Rightarrow \frac{1}{\lambda}=R\left(\frac{1}{n_f^2}-\frac{1}{n_i^2}\right)
\end{aligned}
$
where $R=$ Rydberg's constant $=1.097 \times 10^7 \mathrm{~m}^{-1}$
For Hydrogen-like atoms, the wavelength of an emitted photon during the transition from $\mathrm{n}_{\mathrm{f}}$ orbit to $n_i$ orbit is
$
\frac{1}{\lambda}=R z^2\left(\frac{1}{n_f^2}-\frac{1}{n_i^2}\right)
$

Because of this photon, spectra of hydrogen atoms will emit which are studied by various scientists. One such scientist named Balmer found a formula that gives the wavelengths of these lines for all the transitions taking place to the 2nd orbit.

The Balmer series is a series of spectral emission lines of the hydrogen atom that result from electron transitions from higher levels down to the energy level with principal quantum number 2

Balmer observed the spectra and found the formula for the visible range spectra which is obtained by Balmer's formula is-

$\frac{1}{\lambda}=R\left(\frac{1}{2^2}-\frac{1}{n^2}\right) \ldots$

Here, $n=3,4,5, \ldots \ldots$ etc.
$
R=\text { Rydberg constant }=1.097 \times 10^7 \mathrm{~m}^{-1}
$
and $\lambda$ is the wavelength of the light photon emitted during the transition.

Since Balmer had found the formula for n = 2, we can obtain different spectra for different values of n. For n = $\infty$, we get the smallest wavelength of this series, which is equal to = $3646 \dot{A}$. We can also obtain the value of wavelength for Balmer's series by putting different values of 'n' in the equation (1). Similarly, we can obtain the wavelength of the different spectra like the Lyman, and Paschen series. The Balmer series is in the visible range but the Lyman series is in the Ultraviolet range and the Paschen, Brackett, and Pfund are in the Infrared range.

Lyman series: $\frac{1}{\lambda}=R\left(\frac{1}{1^2}-\frac{1}{n^2}\right), n=2,3,4, \ldots$
Balmer series: $\frac{1}{\lambda}=R\left(\frac{1}{2^2}-\frac{1}{n^2}\right), n=3,4,5, \ldots$
Paschen series: $\frac{1}{\lambda}=R\left(\frac{1}{3^2}-\frac{1}{n^2}\right), n=4,5,6, \ldots$
Brackett series: $\frac{1}{\lambda}=R\left(\frac{1}{4^2}-\frac{1}{n^2}\right), n=5,6,7, \ldots$
Pfund series: $\frac{1}{\lambda}=R\left(\frac{1}{5^2}-\frac{1}{n^2}\right), n=6,7,8$

This is for the hydrogen spectrum

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Solved Examples Based on Line Spectra of the Hydrogen Atom

Example 1: The ratio of the wavelengths of the first line of the Lyman series and the first line of the Balmer series is :

1) 1: 3

2) 27: 5

3) 5: 27

4) 4: 9

Solution:

$
\frac{1}{\lambda}=\mathrm{R}\left(\frac{1}{\mathrm{n}_1^2}-\frac{1}{\mathrm{n}_2^2}\right)
$

For the first line of the Lyman series $\mathrm{n}_1=1$ and $\mathrm{n}_2=2$
For the first line of the Balmer series $\mathrm{n}_2=2$ and $\mathrm{n}_2=3$
So, $\frac{\lambda_{\text {Lymen }}}{\lambda_{\text {Balmer }}}=\frac{5}{27}$

Example 2: The wavelength of the first line of the Balmer series is 6563 Å. The Rydberg constant for hydrogen is about :

1) $1.09 \times 10^7$ per m
2) $1.09 \times 10^8$ per m
3) $1.09 \times 10^9$ per m
4) $1.09 \times 10^5$ per m

Solution:

$\begin{aligned} & \frac{1}{\lambda}=\mathrm{R}\left[\frac{1}{4}-\frac{1}{9}\right]=\frac{5 \mathrm{R}}{36} \\ & \therefore \mathrm{R}=\frac{36}{5 \lambda}=\frac{36}{5 \times 6563 \times 10^{-10}}=1.09 \times 10^7 \mathrm{~m}^{-1}\end{aligned}$

Hence, the answer is the option (1).

Example 3: The ratio of the longest wavelength and the shortest wavelength observed in the five spectral series of the emission spectrum of hydrogen is

1) $\frac{4}{3}$
2) $\frac{525}{376}$
3) 25
4) $\frac{900}{11}$

Solution:

The shortest wavelength comes from $\mathrm{n}_1=\infty$ to $\mathrm{n}_2=1$ and the longest wavelength comes from $\mathrm{n}_1=6$ to $\mathrm{n}_2=5$ in the given case.

Hence
$
\begin{aligned}
& \frac{1}{\lambda_{\min }}=R\left(\frac{1}{1^2}-\frac{1}{\infty^2}\right)=R \\
& \frac{1}{\lambda_{\max }}=R\left(\frac{1}{5^2}-\frac{1}{6^2}\right)=\mathrm{R}\left(\frac{36-25}{25 \times 36}\right)=\frac{11}{900} \mathrm{R} \\
& \therefore \frac{\lambda_{\max }}{\lambda_{\min }}=\frac{900}{11}
\end{aligned}
$

Example 4: The extreme wavelengths of the Paschen series are:

1) $0.365 \mu \mathrm{m}$ and $0.565 \mu \mathrm{m}$
2) $0.818 \mu \mathrm{m}$ and $1.89 \mu \mathrm{m}$
3) $1.45 \mu \mathrm{m}$ and $4.04 \mu \mathrm{m}$
4) $2.27 \mu \mathrm{m}$ and $7.43 \mu \mathrm{m}$

Solution:

$\begin{aligned} & \text { In Paschen series } \frac{1}{\lambda_{\max }}=\mathrm{R}\left[\frac{1}{(3)^2}-\frac{1}{(4)^2}\right] \\ & \Rightarrow \lambda_{\max }=\frac{144}{7 \mathrm{R}}=\frac{144}{7 \times 1.1 \times 10^7}=1.89 \times 10^{-6} \mathrm{~m}=1.89 \mu \mathrm{m} \\ & \text { Similarly } \lambda_{\min }=\frac{9}{\mathrm{R}}=\frac{9}{1.1 \times 10^7}=0.818 \mu \mathrm{m}\end{aligned}$

Hence, the answer is the option (2).

Example 5: The third line of the Balmer series of an ion equivalent to a hydrogen atom has a wavelength of 108.5 mm. The ground state energy of an electron of this ion will be:

1) 3.4 eV
2) 13.6 eV
3) 54.4 eV
4) 122.4 eV

Solution:

For the third line of the Balmer series $\mathrm{n}_1=2, \mathrm{n}_2=5$
$
\therefore \frac{1}{\lambda}=R Z^2\left[\frac{1}{\mathrm{n}_1^2}-\frac{1}{\mathrm{n}_2^2}\right] \text { gives } \mathrm{Z}^2=\frac{\mathrm{n}_1^2 \mathrm{n}_2^2}{\left(\mathrm{n}_2^2-\mathrm{n}_1^2\right) \lambda \mathrm{R}}
$

On putting values $\mathrm{Z}=2$

From $E=-\frac{13.6 \mathrm{Z}^2}{n^2}=\frac{-13.6(2)^2}{(1)^2}=-54.4 \mathrm{eV}$

Hence, the answer is the option (3).

Summary

The line spectra of the hydrogen atom, characterized by discrete emission lines, arise from electrons transitioning between different energy levels. These spectral lines, including those in the Balmer, Lyman, and Paschen series, provide crucial evidence of quantized energy states within the atom. Real-life applications of this knowledge span from identifying chemical compositions in astronomy to practical uses in spectroscopy for material analysis. The study of hydrogen’s line spectra not only enhances our understanding of atomic structure but also drives technological advancements in various scientific fields.

Frequently Asked Questions (FAQs)

Q: How does the concept of oscillator strength relate to the intensities of lines in the hydrogen spectrum?
A:
Oscillator strength is a quantum mechanical quantity that describes the probability of an electron transition between energy levels. In the hydrogen spectrum, the oscillator strength is directly related to the intensity of spectral lines. Transitions with higher oscillator strengths produce stronger, more intense spectral lines. This concept is crucial for predicting and interpreting line intensities in the hydrogen spectrum and helps explain why some lines are brighter than
Q: What is the role of the hydrogen spectrum in the study of quasars and other distant cosmic objects?
A:
The hydrogen spectrum plays a crucial role in studying quasars and other distant cosmic objects. The Lyman-alpha line, in particular, is often used to determine the redshift and thus the distance of these objects. By comparing the observed wavelength of this line to its known rest wavelength, astronomers can calculate how fast the object is moving away from us and estimate its distance. Additionally, absorption lines in the spectra of quasars can reveal information about the intergalactic medium and the distribution of matter between us and the quasar.
Q: How does isotope shift affect the hydrogen spectrum when comparing protium, deuterium, and tritium?
A:
Isotope shift causes slight differences in the spectra of hydrogen isotopes (protium, deuterium, and tritium) due to their different nuclear masses. The heavier nucleus in deuterium and tritium affects the reduced mass of the electron-nucleus system, slightly altering the energy levels. This results in small shifts of spectral lines to higher frequencies (shorter wavelengths) for heavier isotopes. These shifts are small but measurable and are important in high-precision spectroscopy and in studying the abundance of different hydrogen isotopes in various environments.
Q: What is the significance of the Paschen series in the hydrogen spectrum?
A:
The Paschen series is a set of spectral lines in the infrared region of the hydrogen spectrum. These lines result from electron transitions from higher energy levels (n ≥ 4) to the third energy level (n = 3). While less visible than the Balmer series, the Paschen series is important in astrophysics for studying cooler objects like low-mass stars and brown dwarfs. It's also useful in analyzing hydrogen in laboratory plasmas and in understanding the energy level structure of hydrogen-like atoms.
Q: How does the hydrogen spectrum change in high-pressure environments?
A:
In high-pressure environments, the hydrogen spectrum undergoes significant changes. The increased density leads to more frequent collisions between atoms, causing pressure broadening of spectral lines. This broadening can make individual lines less distinct and may cause them to overlap. Additionally, the electric fields from nearby particles can induce Stark broadening. At extremely high pressures, such as those found in the interiors of gas giant planets, the discrete energy levels can even merge into continuous bands, fundamentally altering the spectrum.
Q: What is the Stark effect and how does it affect the hydrogen spectrum?
A:
The Stark effect is the splitting of spectral lines in the presence of an external electric field. In hydrogen, it causes the energy levels to shift and split, resulting in the broadening and splitting of spectral lines. The effect is analogous to the Zeeman effect for magnetic fields. The Stark effect is important in plasma diagnostics and in understanding the spectra of atoms in strong electric fields, such as those near charged particles in plasmas.
Q: What is the relationship between the hydrogen spectrum and the concept of atomic orbitals?
A:
Atomic orbitals represent the probability distributions of electrons around the nucleus. In hydrogen, these orbitals correspond directly to the energy levels that produce the spectral lines. The quantum numbers that define these orbitals (n, l, m) determine the allowed energy states and transitions. The shapes and orientations of these orbitals influence the probabilities of different transitions, which in turn affect the intensities of spectral lines. Understanding atomic orbitals is crucial for interpreting the hydrogen spectrum and predicting transition probabilities.
Q: How does the concept of quantum defect apply to the hydrogen spectrum compared to other elements?
A:
The quantum defect is a measure of how much an atom's energy levels deviate from those predicted by the simple Bohr model. For hydrogen, the quantum defect is essentially zero, meaning its spectrum closely matches the predictions of the Bohr model. However, for other elements, the quantum defect becomes significant due to the screening effect of inner electrons. This concept helps explain why the spectra of other elements differ from hydrogen's, despite following similar principles.
Q: What is the significance of the Inglis-Teller limit in the hydrogen spectrum?
A:
The Inglis-Teller limit is the point at which spectral lines in a series (like the Balmer series) merge into a continuum due to pressure broadening. It's particularly important in dense plasmas and stellar atmospheres. As density increases, electric fields from nearby particles cause Stark broadening of spectral lines. The Inglis-Teller limit helps determine the electron density in plasmas by observing the highest distinguishable spectral line in a series before they merge.
Q: How does the hydrogen spectrum change in extremely strong magnetic fields, such as those found in neutron stars?
A:
In extremely strong magnetic fields, like those in neutron stars, the hydrogen spectrum is dramatically altered. The normal energy level structure breaks down, and the spectrum becomes dominated by Landau levels - quantized orbits of electrons in a magnetic field. This results in a very different set of spectral lines, often evenly spaced and strongly polarized. The study of these spectra is crucial for understanding the physics of neutron stars and other extreme astrophysical environments.