Name: line spectra from heated gases
Uploaded: 2020-09-03T13:38:27.000Z
Duration: 2 min 39 s
Channel: JoVE

K12 Standards

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NC K-12 Science Standard Course of Study

9-12

Physical Science

Matter and its Interactions

PS.PSc.1.4

Emission Spectra

When an atom absorbs energy, the electrons become excited and move to a higher energy level. As the electrons relax down to either a lower energy state or to the ground state, the excess energy is released as a photon.

The wavelength of the absorbed and 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.

An 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. This is the emission spectrum for hydrogen.

The set of spectral lines in the visible light region is known as the Balmer series. It occurs when electrons transition from an energy level higher than n = 3 back down to n = 2. 

The visible light spectrum appears as spectral lines at 410, 434, 486, and 656 nm, which correspond to energy level transitions from n = 3, 4, 5, and 6, respectively, to n = 2. Additional spectral lines can be measured outside of the visible range, such as the Lyman series in the UV region and the Paschen series in the infrared region.

The wavelengths of spectral lines for hydrogen can be predicted using a mathematical expression, where RH is the Rydberg constant, n1 is the principal quantum number of the lower energy level, and n2 is the principal quantum number for the higher energy level. For the Balmer series, n1 = 2.

Since different atoms have different energy levels, the spectral emission lines vary from element to element and are used to identify substances.

The inverse of an emission spectrum is its absorption spectrum. Looking at hydrogen, the lines in its absorption spectrum are located at the same wavelengths of its emission spectrum, but they are dark. These are the wavelengths of light that are absorbed by a hydrogen atom when it is exposed to a continuous white light spectrum.

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 range of energies, and thereby produce a continuous spectrum in which an unbroken series of wavelengths is present.

In contrast to continuous spectra, light can also occur as discrete or line spectra having very narrow linewidths interspersed throughout the spectral regions. Exciting a gas at low partial pressure using an electrical current, or heating it, will produce line spectra. Fluorescent light bulbs and neon signs operate in this way. Each element displays its own characteristic set of lines, as do molecules, although their spectra are generally much more complicated.

Each emission line consists of a single wavelength of light, which implies that the light emitted by a gas consists of a set of discrete energies. For example, when an electric discharge passes through a tube containing hydrogen gas at low pressure, the H2 molecules are broken apart into separate H atoms and a blue-pink color is observed. Passing the light through a prism produces a line spectrum, indicating that this light is composed of photons of four visible wavelengths.

The origin of discrete spectra in atoms and molecules was extremely puzzling to scientists in the late nineteenth century. According to classical electromagnetic theory, only continuous spectra should be observed. Other discrete lines for the hydrogen atom were found in the UV and IR regions. Johannes Rydberg generalized Balmer's work and developed an empirical formula that predicted all of hydrogen's emission lines, not just those restricted to the visible range, where, n1 and n2 are integers, n1 &#60; n2

<img alt="Balmer series equation for hydrogen spectrum, formula illustrating spectral line transitions." src="/files/core/CDNSource/lessons/977/977_Equations-01.png" style="height: 40px;" data-altupdatedat="1761830439000" data-alt="Eq1" />

Even in the late nineteenth century, spectroscopy was a very precise science, and so the wavelengths of hydrogen were measured to very high accuracy, which implied that the Rydberg constant could be determined very precisely as well. That such a simple formula as the Rydberg formula could account for such precise measurements seemed astounding at the time, but it was the eventual explanation for emission spectra by Neils Bohr in 1913 that ultimately convinced scientists to abandon classical physics and spurred the development of modern quantum mechanics.

This text is adapted from <a href="https://openstax.org/books/chemistry-atoms-first-2e/pages/3-1-electromagnetic-energy">Openstax, Chemistry 2e, Section 3.1: Electromagnetic Energy</a>.

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 ...

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KEY TERMS AND CONCEPTS

line spectra from heated gases

Line spectra are produced when heated gases give off light at specific wavelengths. In contrast, hot solids, liquids, or condensed gases radiate excess energy as a continuous spectrum. That spectrum contains an unbroken range of wavelengths.

A gas at low partial pressure can produce a line spectrum when it is excited by an electric current or heated. Fluorescent light bulbs and neon signs work in this way. Each element has its own characteristic set of lines, and molecules can also produce line spectra, although their patterns are usually more complex.

Each emission line is a single wavelength of light. That means the light from a gas is made of discrete energies rather than a smooth spread of values. A good example is hydrogen gas in an electric discharge tube. The H2 molecules break apart into separate H atoms, and the gas gives off a blue-pink color.

When that light is passed through a prism, it separates into a line spectrum. For hydrogen, this shows four visible wavelengths. Scientists also found other hydrogen lines in the ultraviolet and infrared regions. In the late nineteenth century, these discrete spectra were hard to explain because classical electromagnetic theory predicted only continuous spectra.

Johannes Rydberg extended Balmer's work and wrote an empirical formula that predicted all of hydrogen's emission lines, not just the visible ones. The wavelengths of hydrogen were measured with high accuracy, so the Rydberg constant could also be determined very precisely. Later, Niels Bohr explained emission spectra in 1913. His explanation helped push scientists away from classical physics and toward modern quantum mechanics.

Line Spectra from Heated Gases

the wave nature of light

Measuring Light with Wavelength and Frequency

Light can be described as an electromagnetic wave, and its behavior helps explain many patterns seen in science. For centuries, scientists have tried to understand light. Newton used lenses and prisms to show that white light is made of the colors of the rainbow mixed together. He explained light as tiny particles moving at high speed. Other scientists, including Christiaan Huygens, showed that reflection and refraction could also be explained by treating light as a wave. At the time, people...

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the bohr model

Bohr’s Quantized Atomic Orbits

Bohr’s quantized atomic orbits explain why hydrogen atoms can stay stable. Early in the twentieth century, Ernest Rutherford and his colleagues described atoms as tiny dense nuclei with lighter electrons moving around them. This picture was called the planetary model because it compared the atom to a miniature solar system, with electrons orbiting the nucleus like planets around the sun. The simplest atom is hydrogen. It has one proton in the nucleus and one electron moving around it. In the...

Line spectra are produced when heated gases give off light at specific wavelengths. In contrast, hot solids, liquids, or condensed gases radiate excess energy as a continuous spectrum. That spectrum contains an unbroken range of wavelengths. A gas at low partial pressure can produce a line spectrum when it is excited by an electric current or heated. Fluorescent light bulbs and neon signs work in this way. Each element has its own characteristic set of lines, and molecules can also produce...

the de broglie wavelength

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Electron waves and standing orbits explain how tiny particles can behave like waves. Large objects in the macroscopic world, like a billiard ball on a table, follow classical physics. A ball keeps moving in a straight line unless it hits another ball or a force such as friction acts on it. It also has a well-defined position, velocity, and momentum, where momentum is mass times velocity. Waves behave differently from macroscopic particles. When waves interact, they can form interference...

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the quantum-mechanical model of an atom

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Quantum mechanics explains how electrons behave in atoms using waves and probability. Erwin Schrödinger built on de Broglie’s idea that an electron in hydrogen could act like a circular standing wave instead of moving in fixed circular paths. He then derived the Schrödinger equation, which became a major step in modern atomic theory. When Schrödinger applied his equation to hydrogen-like atoms, it matched Bohr’s energy expression and the Rydberg formula for hydrogen spectra. He described...

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the pauli exclusion principle

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Electron spin and orbital pairing explain why only two electrons can fit in one atomic orbital. Electron configuration describes how electrons are arranged in an atom’s orbitals. It is written with three parts: the principal quantum number, n, the orbital type, or subshell, and a superscript that shows how many electrons are in that subshell. For example, 2p4 means four electrons in a p subshell with n = 2, and 3d8 means eight electrons in a d subshell with n = 3. The three quantum numbers...

the energies of atomic orbitals

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Shielding and orbital penetration help explain why electrons in an atom have different energies. Electrons are attracted to the positive nucleus, but they also repel one another in atoms with more than one electron. The strength of these forces depends on the charge and the distance between the particles. Opposite charges attract, like charges repel, and the force gets weaker as the particles move farther apart. An electron that stays closer to the nucleus is held more tightly. That means...

the aufbau principle and hund's rule

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electron configuration of multielectron atoms

electron-configuration-of-multielectron-atoms

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Valence electrons and noble gas shorthand help show how electrons are arranged in multielectron atoms. Sodium, the alkali metal with atomic number 11, has one more electron than neon. That extra electron goes into the lowest-energy subshell available, the 3s orbital, so sodium has the configuration 1s2 2s2 2p6 3s1. Electrons in the outermost shell, where n is highest, are called valence electrons. Electrons in the inner shells are called core electrons. Because core electrons often match a...

key terms and concepts

Explore how heated gases produce line spectra, why hydrogen shows discrete wavelengths, and how Rydberg and Bohr helped explain emission lines.

Line Spectra from Heated Gases | JoVE K12

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Earth and Space Science

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3.3.9-12.B

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Line Spectra from Heated Gases

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