Name: hydrogen and noble gas emission spectra
Uploaded: 2020-03-26T07:22:26.000Z
Duration: 3 min 57 s
Channel: JoVE

K12 Standards

Tennessee Academic Standards

High School

Chemistry

waves-and-their-applications-in-technologies-for-information-transfer

Waves and Their Applications in Technologies for Information Transfer

CHEM1.PS4.1

Balmer Series

<ul><li><a href="#" data-time="0.0">Spectral Emission Lines of Hydrogen, Helium, and Neon
</a></li><li><a href="#" data-time="112">Results
</a></li></ul>

balmer series

<p>Source: Smaa Koraym at Johns Hopkins University, MD, USA</p>
<ol class="note-items">
	<li data-note-item="1" data-parent-collapse="">Spectral Emission Lines of Hydrogen, Helium, and Neon<button class="expand">Expand</button>
	<p>In this experiment, you will use a spectrophotometer to measure the distinct wavelengths of light emitted in the UV and visible range by electron relaxation in hydrogen, helium, and neon. Before starting the lab, make a table in your lab notebook for the elements that you will analyze, the color of the light that you observe from the lamps, and the recorded wavelengths.</p>
	<h4><strong>Table 1: Emission Wavelengths</strong></h4>
	<table border="1" fo:keep-together.within-page="1" fo:keep-with-next.within-page="always">
		<tbody>
			<tr>
				<td style="text-align: center;"><strong>Elements</strong></td>
				<td style="text-align: center;"><strong>Color</strong></td>
				<td style="text-align: center;"><strong>Distinct wavelengths, &#955; (nm)</strong></td>
			</tr>
			<tr>
				<td style="text-align: center;">H</td>
				<td style="text-align: center;"></td>
				<td style="text-align: center;"></td>
			</tr>
			<tr>
				<td style="text-align: center;">He</td>
				<td style="text-align: center;"></td>
				<td style="text-align: center;"></td>
			</tr>
			<tr>
				<td style="text-align: center;">Ne</td>
				<td style="text-align: center;"></td>
				<td style="text-align: center;"></td>
			</tr>
		</tbody>
	</table>
	<p><a href="https://www.jove.com/CDNSource/lm/labs/54/Table_1._Emission_wavelengths.xlsx" target="_blank">Click Here to download Table 1</a></p>
	<ul class="lab-items">
		<li data-sub-note-item="1-2">
		<div class="lab-step">First, put on a lab coat, safety glasses, and nitrile gloves. We suggest students work in pairs for this experiment.</div>
		</li>
		<li data-sub-note-item="1-3">
		<div class="lab-step">Turn on a hand-held spectrophotometer and create a new file for the hydrogen emission spectrum. Configure the spectrophotometer to measure emission intensity with a sampling time of 200 ms, with no sample averaging.</div>
		</li>
		<li data-sub-note-item="1-4">
		<div class="lab-step">When the instructor calls you, bring the spectrophotometer to the hydrogen lamp. The student holding the sensor should move close to the lamp, and the other should be ready to start the spectrophotometer acquisition.</div>
		</li>
		<li data-sub-note-item="1-5">
		<div class="lab-step">When the instructor turns on the lamp, hold the sensor at the center of the lamp and start displaying the spectrum. <strong>Note: </strong>Capture the spectrum as soon as it is clear and has minimal noise, as the lamp cannot be left on for more than 30 s at a time.</div>
		</li>
		<li data-sub-note-item="1-6">
		<div class="lab-step">Save the hydrogen spectrum. Set up another experiment using the same parameters and acquire the spectrum for helium. Acquire a spectrum of neon in the same way.</div>
		</li>
		<li data-sub-note-item="1-7">
		<div class="lab-step">Export your saved data, turn off the spectrophotometer, and put it away.</div>
		</li>
	</ul>
	</li>
	<li data-note-item="2" data-parent-collapse="">Results<button class="expand">Expand</button>
	<ul class="lab-items">
		<li data-sub-note-item="2-1">
		<div class="lab-step">Set up the Rydberg formula to calculate the wavelengths of the Balmer series. <strong>Note:</strong> n<sub>initial</sub> is the number of the energy level where the excited electron starts, and n<sub>final</sub> is the energy level to which the electron relaxes. Set n<sub>final</sub> to 2.
		<h4><strong>Table 2: Frequency and Energy for Each Wavelength</strong></h4>
		<table border="1" fo:keep-together.within-page="1" fo:keep-with-next.within-page="always">
			<tbody>
				<tr>
					<td><strong>Rydberg constant (m<sup>-1</sup>) </strong></td>
					<td style="text-align: center;">1.098 &#215; 10<sup>7</sup></td>
				</tr>
				<tr>
					<td><strong>c (m/s)</strong></td>
					<td style="text-align: center;">2.998 &#215; 10<sup>8</sup></td>
				</tr>
				<tr>
					<td><strong>h (J&#183;s)</strong></td>
					<td style="text-align: center;">6.626 &#215; 10<sup>-34</sup></td>
				</tr>
				<tr>
					<td style="text-align: center;"><strong>ni &#8594; nf</strong></td>
					<td style="text-align: center;"><strong>&#955;<sub>calculated </sub>(nm) </strong></td>
					<td style="text-align: center;"><strong>&#955;<sub>measured</sub> (nm) </strong></td>
					<td style="text-align: center;"><strong>v (THz) </strong></td>
					<td style="text-align: center;"><strong>E (eV) </strong></td>
				</tr>
				<tr>
					<td style="text-align: center;">7 &#8594; 2</td>
					<td></td>
					<td></td>
					<td></td>
					<td></td>
				</tr>
				<tr>
					<td style="text-align: center;">6 &#8594; 2</td>
					<td></td>
					<td></td>
					<td></td>
					<td></td>
				</tr>
				<tr>
					<td style="text-align: center;">5 &#8594; 2</td>
					<td></td>
					<td></td>
					<td></td>
					<td></td>
				</tr>
				<tr>
					<td style="text-align: center;">4 &#8594; 2</td>
					<td></td>
					<td></td>
					<td></td>
					<td></td>
				</tr>
				<tr>
					<td style="text-align: center;">3 &#8594; 2</td>
					<td></td>
					<td></td>
					<td></td>
					<td></td>
				</tr>
			</tbody>
		</table>
		<a href="https://www.jove.com/CDNSource/lm/labs/54/Table_2._Frequency_and_energy_for_each_wavelength.xlsx" target="_blank">Click Here to download Table 2</a></div>
		</li>
		<li data-sub-note-item="2-2">
		<div class="lab-step">Find the wavelength of light emitted by an electron relaxing from level 3 to level 2. Fill in the Rydberg constant for R<sub>H</sub> and solve for 1/&#955;. <strong>Note:</strong> The reciprocal of this value is the wavelength, which can be converted to nm.</div>
		</li>
		<li data-sub-note-item="2-3">
		<div class="lab-step">Calculate the transitions from energy levels 4, 5, 6, and 7 in the same way. <strong>Note: </strong>You should see a good match between your data and your calculations.</div>
		</li>
		<li data-sub-note-item="2-4">
		<div class="lab-step">For each wavelength, calculate the frequency in THz and the energy in eV using the following equations, where c is the speed of light and h is Planck's constant.
		<ul>
			<li><img src="/CDNSource/lm/labs/54/54_Procedure_1.jpg" alt="Frequency equation \( \nu = \frac{c}{\lambda} \) in wave speed calculation formula." data-altupdatedat="1761830439000" /></li>
			<li><img src="/CDNSource/lm/labs/54/54_Procedure_2.jpg" alt="Photon energy formula \(E=h\nu\), equation for calculating energy in quantum mechanics studies." data-altupdatedat="1761830439000" /></li>
		</ul>
		</div>
		</li>
	</ul>
	<p>Look at the calculated energies. You should see a trend of the energy gap increasing by less with each added energy level. This is reflected in the spacing of the peaks in the hydrogen emission spectrum. Helium and neon share this energy level spacing trend.</p>
	<p>From the helium spectrum, adding even one more electron makes the spectral series harder to calculate and identify. This is even more apparent in neon, which has eight more electrons than helium.</p>
	</li>
</ol>

Source: Smaa Koraym at Johns Hopkins University, MD, USA

	Spectral Emission Lines of Hydrogen, Helium, and NeonExpand
	In this experiment, you will use a ...

Education

JoVE Lab Manual

Lab Manual Sub Articles

hydrogen and noble gas emission spectra

<p class="jove_content">Hydrogen and noble gas emission spectra show the light emitted when excited electrons relax to lower energy levels. In this lab, students use a spectrophotometer to measure distinct wavelengths in the UV and visible range from hydrogen, helium, and neon lamps. Before starting, students prepare a table in their lab notebook for the element, the color they observe, and the wavelengths they record.</p>

<p class="jove_content">Students wear a lab coat, safety glasses, and nitrile gloves, and the work is best done in pairs. They turn on a hand-held spectrophotometer and create a new file for the hydrogen emission spectrum. The instrument is set to measure emission intensity with a sampling time of 200 ms and no sample averaging. When the instructor turns on the lamp, students place the sensor at the center of the lamp, start the acquisition, and save the spectrum as soon as it is clear and has minimal noise.</p>

<p class="jove_content">The same setup is then used for helium and neon. Students export the saved data, turn off the spectrophotometer, and put it away. In the results, they use the Rydberg formula to calculate the wavelengths of the Balmer series, which comes from electron transitions to energy level 2. They calculate the transitions from levels 3, 4, 5, 6, and 7 to 2 and compare the calculated wavelengths with the measured values.</p>

<p class="jove_content">Table 2 also includes frequency in THz and energy in eV for each wavelength. Students use the speed of light and Planck’s constant to find these values. The calculated energies show a trend in which the energy gap increases by less with each added energy level, and that pattern matches the spacing of peaks in the hydrogen emission spectrum. Helium and neon show the same spacing trend, but adding even one more electron makes the spectral series harder to calculate and identify. This is even more noticeable in neon, which has eight more electrons than helium.</p>

Hydrogen and Noble Gas Emission Spectra

Gas Discharge Lamp Lab Setup

Gas discharge lamp lab setup prepares a class to collect spectra from hydrogen, helium, and neon. The setup is designed for 10 students working in pairs, with some excess materials. Quantities can be adjusted as needed for the lab section.

Hydrogen and noble gas emission spectra show the light emitted when excited electrons relax to lower energy levels. In this lab, students use a spectrophotometer to measure distinct wavelengths in the UV and visible range from hydrogen, helium, and neon lamps. Before starting, students prepare a table in their lab notebook for the element, the color they observe, and the wavelengths they record.<br /><br />Students wear a lab coat, safety glasses, and nitrile gloves, and the work is best done in pairs.

lab manual sub articles

Measure hydrogen, helium, and neon emission spectra with a spectrophotometer, then calculate Balmer series wavelengths, frequency, and energy.

Hydrogen and Noble Gas Emission Spectra | JoVE K12

<p>Atoms and molecules exhibit very interesting behavior when they absorb and release energy. The electrons in an atom traditionally exist at their lowest energy state, called the ground state, and labeled as n equal to one. However, when an atom absorbs energy, the electrons become excited and move to a higher energy level. When the electrons relax down to either a lower energy state or to the ground state, the excess energy is released as emitted light.</p>
<p>You may be familiar with these energy states, as they are depicted in the Bohr Model, which describes an atom as a nucleus with orbiting electrons in shells or orbitals. These shells are the same as the energy levels and are labeled by n. The wavelength of the emitted light depends on the difference between the high and low energy levels. High energy emitted light results from electrons relaxing from a higher energy level, and low energy emitted light results from electrons relaxing from a lower energy level.</p>
<p>The emission spectrum is a measure of emitted radiation across a range of wavelengths. With pure elemental species, the emission behavior appears as lines of specific wavelengths rather than a broad spectrum. Since different ato</p>

hydrogen emission lines

<p class="jove_content">Hydrogen emission lines show how electrons move between fixed energy levels in the Bohr model. Niels Bohr proposed this model for the hydrogen atom in 1913. In this model, an electron can only occupy certain orbits, and it cannot exist between them.</p>

<p class="jove_content">These energy states are labeled with the quantum number n. The ground state is n = 1. Higher values, such as n = 2, 3, 4, and so on, are excited states. An electron changes state only by absorbing or emitting energy equal to the difference between two levels. A photon with the right energy can move an electron from one orbit to another.</p>

<p class="jove_content">Bohr’s model also links each energy level to a specific amount of potential energy. That value can be calculated with an equation that uses the Rydberg constant, Planck’s constant, and the speed of light. When an electron returns to the ground state or another lower excited state, it releases the extra energy as a photon. The photon energy matches the gap between the two states, so it also determines the wavelength of the emitted light.</p>

<p class="jove_content">Because each atom has its own energy levels, each element gives off its own pattern of light. A mixed sample can produce a continuous spectrum with many wavelengths. A pure sample of one element produces only certain wavelengths. When that light passes through a prism, it appears as separate lines instead of one smooth band.</p>

<p class="jove_content">Hydrogen shows this pattern clearly. Exciting hydrogen atoms requires enough energy to split the H2 molecule, and the extra energy excites the hydrogen electrons to higher levels. When those electrons drop back down, they emit light at wavelengths set by the energy difference between the initial level, n i, and the final level, n f. In pure hydrogen gas, this emission spectrum contains distinct visible, ultraviolet, and infrared lines that are unique to hydrogen.</p>

<p class="jove_content">The visible lines in hydrogen are called the Balmer series. They occur when an electron falls to n = 2 from a higher level. Johann Balmer observed lines at 410.2 nm, 434.1 nm, 486.1 nm, and 656.3 nm, which match transitions from n = 6, 5, 4, and 3 to n = 2. He described these wavelengths with the Balmer formula, where λ is the observed wavelength, C is a constant of 364.50682 nm, n is the lower level, and m is the higher level.</p>

<p class="jove_content">Later, Johannes Rydberg refined this idea using the Rydberg constant. For the Balmer series, the final level n f equals 2. The same relationship, combined with the Bohr model, can be used to calculate the energy change, ΔE, for electron transitions. Other hydrogen spectral series were later identified as well, including the Lyman series, which lies in the ultraviolet region.</p>

Hydrogen Emission Lines

The Bohr Model
Niels Bohr proposed a model for the hydrogen atom in 1913 that described discrete energy states are associated with a fixed electron orbit ...

Test your understanding of Balmer Series with this Chemistry assessment quiz on JoVE Lab Manual. Covers key concepts through video-based questions designed to evaluate student learning.

Balmer Series in Chemistry - Assessment Quiz | JoVE Lab Manual

Prepare a high school spectroscopy lab with hydrogen, helium, and neon discharge lamps, UV-Vis spectrophotometers, and safe handling steps.

Gas Discharge Lamp Lab Setup | JoVE K12

Emission Spectra

When solids, liquids, or condensed gases are heated sufficiently, they radiate some of the excess energy as light. Photons produced in this manner have a ...

The Bohr Model

Following the work of Ernest Rutherford and his colleagues in the early twentieth century, the picture of atoms consisting of tiny dense nuclei surrounded ...

Photoelectric Effect

When light of a particular wavelength strikes a metal surface, electrons are emitted. This is called the photoelectric effect. The minimum frequency of ...

The Wave Nature of Light

The nature of light has been a subject of inquiry since antiquity. In the seventeenth century, Isaac Newton performed experiments with lenses and prisms ...

dual-nature-of-electromagnetic-em-radiation

Dual Nature of Electromagnetic (EM) Radiation

Electromagnetic (EM) radiation consists of electric and magnetic field components oscillating in planes perpendicular to each other and mutually ...

Source: Smaa Koraym at Johns Hopkins University, MD, USA

	Preparation of the LaboratoryExpand
	Here, we show the laboratory preparation for 10 students ...

The de Broglie Wavelength

In the macroscopic world, objects that are large enough to be seen by the naked eye follow the rules of classical physics. A billiard ball moving on a ...

Generating Electromagnetic Radiations

The German physicist Heinrich Hertz (1857&#8211;1894) was the first to generate and detect certain types of electromagnetic waves in the laboratory. ...

Using a model, explain why elements emit characteristic frequencies of light and how this information is used. EXPLANATION: An understanding of the behavior of light as has been developed first by investigating properties of waves in fourth grade and more recently in eighth grade when students differentiated between mechanical and electromagnetic waves. This is the first time that students investigate a mechanism for the emission of photons. Emphasis should be placed on the emission of characteristic colors of light emitted when electrons undergo specific movements and unique spectra of each element that result. Students should explore the implications of these ideas on astronomy. (Clarification may be needed to differentiate this particulate behavior of light from previous discussions of its wave properties; however, discussions of Quantum Theory in differentiating wave-particle duality are beyond the scope of this standard)

Hydrogen and Noble Gas Emission Spectra

Procedure

Transcript

Tags

More videos by Standards