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

Quand un atome absorbe l'énergie, les électrons deviennent excités et se déplacent à un niveau d'énergie plus élevé. Lorsque les électrons se détendent vers le bas à un état d'énergie plus faible ou à l'état fondamental, l'énergie excédentaire est libérée sous forme de photon. La longueur d'onde de la lumière absorbée et émise dépend de la différence entre les niveaux d'énergie haute et basse.La lumière émise à haute énergie résulte d'électrons se relaxant à partir d'un niveau d'énergie plus élevé, et la lumière émise à faible énergie résulte d'électrons se relaxant à partir d'un niveau d'énergie plus faible. Un spectre d'émission est une mesure du rayonnement émis à travers une gamme de longueurs d'onde. Avec les espèces élémentaires pures, le comportement d'émission apparaît comme des lignes de longueurs d'onde spécifiques plutôt qu'un large spectre.C'est le spectre d'émission de l'hydrogène. L'ensemble des raies spectrales dans la région de la lumière visible est connu comme la série Balmer. Il se produit lorsque les électrons passent d'un niveau d'énergie supérieur à n 3 de retour à n 2.Le spectre de la lumière visible apparaît sous forme de raies spectrales à 410, 434, 486 et 656 nm, qui correspondent aux transitions de niveau d'énergie de n 3, 4, 5 et 6, respectivement, à n 2. Des raies spectrales supplémentaires peuvent être mesurées en dehors de la plage visible, telles que la série Lyman dans la région UV et la série Paschen dans la région infrarouge. Les longueurs d'onde des raies spectrales pour l'hydrogène peuvent être prédites en utilisant une expression mathématique, où R-H est la constante de Rydberg, n1 est le nombre quantique principal du niveau d'énergie inférieur, et n2 est le nombre quantique principal du niveau d'énergie supérieur.Pour la série Balmer, n1 2. Puisque différents atomes ont des niveaux d'énergie différents, les raies d'émission spectrales varient d'un élément à l'autre et sont utilisées pour identifier les substances. L'inverse d'un spectre d'émission est son spectre d'absorption.En regardant l'hydrogène, les raies dans son spectre d'absorption sont situées aux mêmes longueurs d'onde de son spectre d'émission, mais elles sont sombres. Ce sont les longueurs d'onde de la lumière qui sont absorbées par un atome d'hydrogène lorsqu'il est exposé à un spectre de lumière blanche continue.

spectres d'émission

Lorsque les solides, les liquides ou les gaz condensés sont suffisamment chauffés, ils émettent une partie de l'énergie excédentaire sous forme de lumière. Les photons produits de cette manière se situent sur une gamme d'énergies, et produisent ainsi un spectre continu dans lequel se trouve une série ininterrompue de longueurs d'onde. 
 
 Contrairement aux spectres continus, la lumière peut également se produire sous forme de spectres discontinus ou de raies de largeurs très fines répandues le long des régions du spectre. L'excitation d'un gaz avec une pression partielle basse à l'aide d'un courant électrique, ou son chauffage, produira des spectres de raies. Les ampoules fluorescentes et les enseignes au néon fonctionnent de cette façon. Chaque élément affiche son propre ensemble de raies caractéristique, tout comme les molécules, bien que leurs spectres soient généralement beaucoup plus compliqués. 
 
 Chaque raie d'émission consiste en une seule longueur d'onde de lumière, ce qui implique que la lumière émise par un gaz est constituée d'un ensemble d'énergies distinctes. Par exemple, lorsqu'une décharge électrique passe à travers un tube contenant de l'hydrogène gazeux à basse pression, les molécules de H2 sont séparées en atomes H séparés et une couleur bleu-rose est observée. Le passage de la lumière à travers un prisme produit un spectre de raies, indiquant que cette lumière est composée de photons de quatre longueurs d'onde visibles. 
 
 L'origine des spectres discontinus dans les atomes et les molécules était extrêmement troublante pour les scientifiques à la fin du XIXe siècle. Selon la théorie électromagnétique classique, seuls les spectres continus doivent être observés. D'autres raies distinctes pour l'atome d'hydrogène ont été trouvées dans les régions UV et IR. Johannes Rydberg généralisa le travail de Balmer et développa une formule empirique qui prédisait toutes les raies d'émission de l'hydrogène, pas seulement celles limitées au visible, où, n1 et n2 étant des nombres entiers, n1 < n2
 
 <img alt="Eq1" src="https://www.jovecore.com/CDNSource/lessons/977/977_Equations-01.png" style="height: 40px;" />
 
 Même à la fin du XIXe siècle, la spectroscopie était une science très précise, et donc les longueurs d'onde de l'hydrogène ont été mesurées avec une très grande exactitude, ce qui implique que la constante de Rydberg pouvait aussi être déterminée de façon très précise. Qu'une formule aussi simple que celle de Rydberg puisse expliquer des mesures aussi précises semblait stupéfiant à l'époque, mais c'est l'explication finale des spectres d'émission par Neils Bohr en 1913 qui a finalement convaincu les scientifiques d'abandonner la physique classique et a stimulé le développement de la mécanique quantique moderne. 
 
 Ce texte est adapté de <a href="https://openstax.org/books/chemistry-atoms-first-2e/pages/3-1-electromagnetic-energy">Openstax, Chimie 2e, Section 3.1 : Énergie électromagnétique</a>.

Spectres d'émission

Lorsque les solides, les liquides ou les gaz condensés sont suffisamment chauffés, ils émettent une partie de l’énergie excédentaire sous forme de lumière. Les photons produits de cette manière ont un ...

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Electronic Structure of Atoms

Structure électronique des atomes

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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...

the electromagnetic spectrum

Electromagnetic Waves: Uses Across the Spectrum

The electromagnetic spectrum includes all electromagnetic radiation arranged by frequency and wavelength. Visible light is only a small part of it, and each color of visible light has its own frequency and wavelength. Different parts of the spectrum also use different units. Radio waves are usually measured in MHz, while the visible range is often measured in nm or angstroms. The spectrum is organized from radio waves to gamma rays as frequency increases and wavelength decreases. Radio waves...

interference and diffraction

Wave Patterns from Light and Slits

Light can form wave patterns through interference and diffraction. When two electromagnetic waves meet with their peaks and troughs lined up, they combine into a larger wave. This is constructive interference, and the waves are in phase. When a peak of one wave meets a trough of the other, the waves reduce each other. This is destructive interference, and the result has a much lower amplitude. Diffraction is another wave behavior. When light passes through a slit that is about the same size as...

photoelectric effect

Why Light Ejects Electrons from Metals

Light can eject electrons from a metal surface when it has the right frequency. This process is called the photoelectric effect. Each metal has its own threshold frequency, which is the minimum light frequency needed to free electrons. Light below the threshold frequency cannot cause emission, even if it is very bright. When the light frequency is above the threshold, electrons are ejected. In that case, the number of electrons released is directly proportional to the intensity, or brightness,...

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...

emission spectra

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

Electron Waves and Standing Orbits

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...

the uncertainty principle

Electron Position and Momentum Limits

Heisenberg’s uncertainty principle sets a limit on how precisely we can know an electron’s position and momentum at the same time. Momentum is the motion of a particle, and the two measurements affect each other. If the momentum is known more accurately, the position becomes less certain, and the reverse is also true. Heisenberg described this limit with an equation that uses Planck’s constant. The equation shows the smallest possible uncertainty for position and momentum together. It explains...

the quantum-mechanical model of an atom

Electron Orbitals in Quantum Mechanics

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...

quantum numbers

Electron Orbitals and Spin

Electron orbitals are described by quantum numbers that show where an electron is most likely to be found around the nucleus. These numbers come from the quantum mechanical model, which is based on solutions of the Schrödinger equation. They also help explain why electron energy comes in set values instead of changing smoothly. The principal quantum number, n, labels the energy level. It can be 1, 2, 3, and so on, and higher n values usually mean higher electron energy. This number also tells...

atomic orbitals

Electron Regions in Atomic Orbitals

Electron regions in atomic orbitals show where an electron is most likely to be found in an atom. These three-dimensional regions are often drawn as boundary surfaces around the densest part of the cloud. The shape of an orbital depends on the angular momentum quantum number, l. This value can range from 0 to n - 1, where n is the principal quantum number. For n = 1, only l = 0 is allowed. For n = 2, both l = 0 and l = 1 are allowed. Orbitals with the same value of l form a subshell. The l =...

the pauli exclusion principle

Electron Spin and Orbital Pairing

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

Shielding and Orbital Penetration

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

Electron Filling Order in Atomic Orbitals

Electron filling order in atomic orbitals follows the aufbau principle and Hund's rule. To write an electron configuration, we build atoms in order of atomic number. Starting with hydrogen and moving across the periods of the periodic table, one proton is added to the nucleus and one electron is placed into the correct subshell. This process describes the ground-state electron configurations of the elements. The aufbau principle comes from the German word aufbau, which means "to build up." It...

electron configuration of multielectron atoms

electron-configuration-of-multielectron-atoms

Valence Electrons and Noble Gas Shorthand

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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3.3.9-12.B

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Earth’s Place In The Universe

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Waves and Their Applications in Technologies for Information Transfer

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Matter and Its Interactions

9-12.C.1.1.c

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Light Emission and Electron Movement

Classifying Matter by State

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Classifying Matter by Composition

Matter: Pure Substances and Mixtures
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PS.PSc.1.1

Phase Transitions

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Phase Transitions: Vaporization and Condensation

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Vapor Pressure

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Clausius-Clapeyron Equation

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Phase Transitions: Melting and Freezing

Heating a crystalline solid increases the average energy of its atoms, molecules, or ions, and the solid gets hotter. At some point, the added energy ...

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Phase Transitions: Sublimation and Deposition

Some solids can transition directly into the gaseous state, bypassing the liquid state, via a process known as sublimation. At room temperature and ...

Heating and Cooling Curves

When a substance&#8212;isolated from its environment&#8212;is subjected to heat changes, corresponding changes in temperature and phase of the substance ...

Phase Diagrams

A phase diagram combines plots of pressure versus temperature for the liquid-gas, solid-liquid, and solid-gas phase-transition equilibria of a substance. ...

Use models to compare the phases of matter and the physical changes they undergo.

PS.PSc.1.2

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Physical and Chemical Properties of Matter

The characteristics that enable us to distinguish one substance from another are called properties.

Physical Properties of Matter

A physical property is ...

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Molecular Comparison of Gases, Liquids, and Solids

Particles in a solid are tightly packed together (fixed shape) and often arranged in a regular pattern; in a liquid, they are close together with no ...

Intermolecular vs Intramolecular Forces

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Intermolecular Forces

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Comparing Intermolecular Forces: Melting Point, Boiling Point, and Miscibility

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PS.PSc.1.3

The Atomic Theory of Matter

The earliest recorded discussion of the basic structure of matter comes from ancient Greek philosophers. Leucippus and Democritus argued that all matter ...

Subatomic Particles

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Understand types, properties, and structure of matter.

Line Spectra from Heated Gases

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