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

K12 Standards

north-carolina-k-12-science-standard-course-of-study

NC K-12 Science Standard Course of Study

9-12

Physical Science

Matter and its Interactions

PS.PSc.1.4

Emission Spectra

Wenn ein Atom Energie aufnimmt, werden die Elektronen angeregt und bewegen sich auf ein höheres Energieniveau zu. Während sich die Elektronen entweder in einen niedrigeren Energiezustand oder in den Grundzustand entspannen, wird die überschüssige Energie als Photon freigesetzt. Die Wellenlänge des absorbierten und emittierten Lichts hängt von der Differenz zwischen des höheren und des niedrigeren Energiezustandes ab.Das von der höheren Energie emittierte Licht ist Ergebnis von Elektronen, die sich in einem höheren Energiezustand befinden, und das von der niedrigeren Energie emittierte Licht ist Ergebnis von Elektronen, die sich in einem niedrigeren Energiezustand befinden. Ein Emissionsspektrum ist ein Maß der emittierten Strahlung, über einen Bereich von Wellenlängen. Bei reinen Elementen erscheint das Emissionsverhalten eher als Linien bestimmter Wellenlängen und nicht als ein breites Spektrum.Das ist das Emissionsspektrum für Wasserstoff. Die Menge der Spektrallinien im Bereich des sichtbaren Lichts ist als die Balmer-Serie bekannt. Es tritt auf, wenn Elektronen von einem Energiezustand höher als n 3 zurück auf n 2 übergehen.Das sichtbare Lichtspektrum erscheint als Spektrallinien bei 410, 434, 486, und 656 nm, die den Energieübergängen von n 3, 4, 5 bzw. 6 bis n 2 entspricht. Zusätzliche Spektrallinien können außerhalb des sichtbaren Bereichs gemessen werden, wie zum Beispiel der Lyman-Serie im UV-Bereich und die Paschen-Serie im Infrarotbereich.Die Wellenlängen der Spektrallinien für Wasserstoff können mit Hilfe einer mathematischen Formel vorhergesagt werden, wobei R-H die Rydberg-Konstante ist, n1 ist die Hauptquantenzahl des unteren Energiezustandes und n2 ist die Hauptquantenzahl für den höheren Energiezustand. Für die Balmer-Reihe gilt n1 2. Da verschiedene Atome unterschiedliche Energiezustände haben, variieren die spektralen Emissionslinien von Element zu Element und werden zur Identifizierung von Stoffen verwendet.Der Kehrwert eines Emissionsspektrums ist sein Absorptionsspektrum. Betrachtet man Wasserstoff, so befinden sich die Linien in seinem Absorptionsspektrum auf den gleichen Wellenlängen wie in seinem Emissionspektrums, aber sie sind dunkel. Das sind die Wellenlängen des Lichts, die von einem Wasserstoffatom absorbiert werden, wenn es einem kontinuierlichen weißen Lichtspektrum ausgesetzt ist.

emissionsspektren

Wenn Feststoffe, Flüssigkeiten oder kondensierte Gase ausreichend erhitzt werden, strahlen sie einen Teil der überschüssigen Energie als Licht ab. Auf diese Weise erzeugte Photonen haben unterschiedliche Energien und erzeugen dadurch ein kontinuierliches Spektrum, in dem eine ununterbrochene Folge von Wellenlängen vorhanden ist. Im Gegensatz zu kontinuierlichen Spektren kann Licht auch als diskrete Spektren oder Linienspektren mit sehr schmalen Bandbreiten auftreten, welche über die Spektralbereiche verteilt sind. Wenn man ein Gas bei niedrigem Partialdruck mit elektrischem Strom anregt oder es erhitzt, entstehen beispielsweise Linienspektren. Auf diese Weise funktionieren Leuchtstofflampen oder Leuchtreklamen. Jedes Element weist seine eigene charakteristische Reihe von Linien auf, das selbe gilt für Moleküle, obwohl ihre Spektren im Allgemeinen wesentlich komplizierter sind. Jede Emissionslinie besteht aus einer einzelnen Lichtwellenlänge, was bedeutet, dass das von einem Gas emittierte Licht aus einer Reihe diskreter Energien besteht. Wenn beispielsweise eine elektrische Entladung durch ein Rohr mit Wasserstoffgas bei niedrigem Druck fließt, werden die H2 -Moleküle in einzelne H-Atome zerlegt und es wird eine blau-rosa Farbe beobachtet. Wenn das Licht durch ein Prisma geleitet wird, erhält man ein Linienspektrum, welches darlegt, dass das Licht aus Photonen mit vier unterschiedlichen Wellenlängen besteht. Der Ursprung diskreter Spektren in Atomen und Molekülen war für Wissenschaftler im späten 19. Jahrhundert äußerst rätselhaft. Nach der klassischen elektromagnetischen Theorie sollten nur kontinuierliche Spektren beobachtet werden. Weitere diskrete Linien für das Wasserstoffatom wurden im UV- und IR-Bereich gefunden. Johannes Rydberg verallgemeinerte Balmers Arbeit und entwickelte eine empirische Formel, welche alle Emissionslinien von Wasserstoff vorhersagt, nicht nur diejenigen, die auf den sichtbaren Bereich beschränkt waren. Hierbei sind n1 und n2 ganze Zahlen sowie n1 < n2 <img alt="Eq1" src="https://www.jove.com/files/core/CDNSource/lessons/977/977_Equations-01.png" style="height: 40px;"> Schon im späten 19. Jahrhundert war die Spektroskopie eine vergleichsweise präzise Wissenschaft. So konnten die Wellenlängen von Wasserstoff mit sehr hoher Genauigkeit gemessen und somit auch die Rydberg-Konstante mit einer hohen Bestimmtheit ermittelt werden. Dass eine so einfache Formel wie die Rydberg-Formel solch präzise Messungen erklären konnte, schien damals erstaunlich, doch war es letztendlich die Erklärung für Emissionsspektren durch Neils Bohr im Jahr 1913, welche die Wissenschaftler davon überzeugte, die klassische Physik aufzugeben und die Entwicklung moderner Quantenphysik voranzutreiben. Dieser Text wurde angepasst von <a href="https://openstax.org/books/chemistry-atoms-first-2e/pages/3-1-electromagnetic-energy">Openstax, Chemistry 2e, Section 3.1: Electromagnetic Energy</a>.

Emissionsspektren

Wenn Festkörper, Flüssigkeiten oder kondensierte Gase ausreichend erhitzt werden, strahlen sie einen Teil der überschüssigen Energie in Form von Licht ab. Photonen, die auf diese Weise erzeugt werden, haben eine ...

Education

Lehre

JoVE Core

Chemistry

Chemie

Eins

Electronic Structure of Atoms

Elektronische Struktur von Atomen

KEY TERMS AND CONCEPTS

KERNBEGRIFFE UND KONZEPTE

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

chemistry

JoVE Core Chemie erweckt allgemeine Konzepte der Chemie durch kompakte und leicht verständliche animierte Videolektionen zum Leben. Außerdem zeigen wir in den Videos Wissenschaftler in Aktion bei klassischen und neuartigen Forschungsexperimenten, wie sie heutzutage in Laboren weltweit durchgeführt wurden.

Pennsylvania State Academic Standards

Earth and Space Science

Space Systems

3.3.9-12.B

3.3.9-12.C

Louisiana Student Standards for Science

High School

Earth Science

Earth’s Place In The Universe

HS-ESS1-1

HS-ESS1-2

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Motion And Stability: Forces And Interactions

HS-PS2-6

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Ohio’s Learning Standards and Model Curriculum

Study Of Matter

Atoms

Tennessee Academic Standards

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Waves and Their Applications in Technologies for Information Transfer

CHEM1.PS4.1

Virginia  Standards of Learning

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CH.2

CH.2.e

PS.2

PS.2.a

Arizona Science Standards

HS+C.P1

HS+C.P1U1.2

Texas Essential Knowledge and Skills

CHEM.6

CHEM.6.A

CHEM.6.B

CHEM.6.C

Alabama Course of Study Science

Matter and Its Interactions

9-12.C.1.1.c

Earth's Place in the Universe

9-12.ESS.1.1.b

9-12.ESS.1.2

9-12.ESS.1.2.c

9-12.ESS.1.3.a

Next Generation Science Standards

Earth’s Place in the Universe

Florida Department of Education

SC.912.E.5

SC.912.E.5.8

SC.912.P.8

SC.912.P.8.3

SC.912.P.10

SC.912.P.10.9

Physics

PHYS.8

PHYS.8.F

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PHYS.9.A

Chemistry and Physics

IPC.7

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South Carolina College and Career Ready Science Standards

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New York State Science Learning Standards

HS. Chemical Reactions

HS-PS1-7.

Georgia Standards of Excellence

SC1.

Light Emission and Electron Movement

Classifying Matter by State

Chemistry is the study of matter and the changes it undergoes. Matter is anything that has mass and occupies space. Matter is all around us; the air, ...

Classifying Matter by Composition

Matter: Pure Substances and Mixtures
According to its composition, the matter can be classified into two broad categories &mdash; pure substances and ...

Elements: Chemical Symbols and Isotopes

A chemical symbol is an abbreviation used to indicate an element or an atom of an element. For example, the symbol for mercury is Hg. The same symbol is ...

Molecules and Compounds

Atoms and Molecules

Everything in the universe is made up of matter, and matter is composed of a combination of elements. An atom is the smallest unit of ...

Solution Formation

There is no one solvent that can dissolve every type of solute. Some substances that readily dissolve in a certain solvent might be insoluble in a ...

Colloids

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Construct an explanation to classify matter as a pure substance or mixture; homogeneous or heterogeneous; element or compound; solution, colloid or suspension.

PS.PSc.1.1

Phase Transitions

Whether solid, liquid, or gas, a substance&#39;s state depends on the order and arrangement of its particles (atoms, molecules, or ions). Particles in the ...

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Phase Transitions: Vaporization and Condensation

The physical form of a substance changes on changing its temperature. For example, raising the temperature of a liquid causes the liquid to vaporize ...

Vapor Pressure

When a liquid vaporizes in a closed container, gas molecules cannot escape. As these gas phase molecules move randomly about, they will occasionally ...

Clausius-Clapeyron Equation

The equilibrium between a liquid and its vapor depends on the temperature of the system; a rise in temperature causes a corresponding rise in the vapor ...

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 ...

phase-transitions-sublimation-and-deposition

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

Intermolecular forces (IMF) are electrostatic attractions arising from charge-charge interactions between molecules. The strength of the intermolecular ...

Intermolecular Forces

Atoms and molecules interact through bonds (or forces): intramolecular and intermolecular. The forces are electrostatic as they arise from interactions ...

comparing-intermolecular-forces-melting-point-boiling-point

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

Intermolecular forces are attractive forces that exist between molecules. They dictate several bulk properties, such as melting points, boiling points, ...

Carry out investigations to compare physical and chemical properties of matter.

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

Dalton was only partially correct about the particles that make up matter. All matter is composed of atoms, and atoms are composed of three smaller ...

Use models to interpret the data presented in Bohr diagrams and electron dot diagrams for neutral atoms of elements 1 through 18.

Understand types, properties, and structure of matter.

Line Spectra from Heated Gases

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