Name: hydrogen emission lines
Uploaded: 2020-03-26T07:22:25.000Z
Duration: 3 min 53 s
Channel: JoVE

K12 Standards

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Chemistry and Physics

IPC.7

IPC.7.E

Balmer Series

Atome und Moleküle zeigen ein sehr interessantes Verhalten, wenn sie Energie absorbieren und wieder abgeben. Die Elektronen in einem Atom existieren traditionell in ihrem niedrigsten Energiezustand, dem sogenannten Grundzustand, und werden als n gleich eins bezeichnet. Wenn ein Atom jedoch Energie absorbiert, werden die Elektronen angeregt und bewegen sich auf ein höheres Energieniveau. Wenn sich die Elektronen entweder in einen niedrigeren Energiezustand oder in den Grundzustand entspannen, wird die überschüssige Energie als emittiertes Licht freigesetzt.
Vielleicht kennen Sie diese Energiezustände, wie sie im Bohr-Modell dargestellt sind, das ein Atom als Kern mit umkreisenden Elektronen in Schalen oder Orbitalen beschreibt. Diese Schalen sind die gleichen wie die Energieniveaus und mit n gekennzeichnet. Die Wellenlänge des emittierten Lichts hängt von der Differenz zwischen dem hohen und dem niedrigen Energieniveau ab. Emittiertes Licht mit hoher Energie resultiert aus Elektronen, die sich von einem höheren Energieniveau entspannen, und emittiertes Licht mit niedriger Energie resultiert aus Elektronen, die sich von einem niedrigeren Energieniveau entspannen.
Das Emissionsspektrum ist ein Maß für die emittierte Strahlung über einen Bereich von Wellenlängen. Bei reinen Elementspezies erscheint das Emissionsverhalten als Linien spezifischer Wellenlängen und nicht als breites Spektrum. Da verschiedene Atome unterschiedliche Energieniveaus haben, variieren diese Spektrallinien von Element zu Element und hängen von den Übergängen ab, die diese Elektronen zwischen den Energiezuständen machen, wenn sie angeregt werden. Zum Beispiel gibt es sechs benannte Reihen von Spektrallinien für Wasserstoff, von denen eine die Balmer-Reihe ist.
Die Balmer-Reihe der Spektrallinien tritt auf, wenn Elektronen von einem Energieniveau über n = 3 zurück zu n = 2 übergehen. Das sichtbare Lichtspektrum für die Balmer-Serie erscheint als Spektrallinien bei 410, 434, 486 und 656 nm. Die h-Alpha-Linie ist die rote Linie bei 656 nm und entsteht durch den Übergang von n= 3 zu n = 2. Die cyanfarbenen, blauen und violetten Linien entsprechen den Übergängen des Energieniveaus von n = 4, 5 bzw. 6 zurück zu n = 2. Zusätzliche Spektrallinien können außerhalb des sichtbaren Bereichs gemessen werden.
Johann Balmer quantifizierte die sichtbaren Linien mit der Balmer-Formel. Dabei ist Lambda die beobachtete Wellenlänge, c ist eine Konstante, n ist das untere Energieniveau von zwei und m ist das höhere Energieniveau. Die Kombination aus der Balmer-Gleichung und dem Bohr-Modell ergibt die Rydberg-Gleichung, die die Spektrallinien vieler verschiedener Elemente beschreibt. In dieser Gleichung ist Lambda die aufgezeichnete Wellenlänge und RH die Rydberg-Konstante. Die Anfangsniveaus, die durch n-initial gekennzeichnet sind, stellen das höhere Energieniveau dar, auf das die Elektronen anregen, und n final ist das niedrigere Energieniveau, auf das sich die Elektronen zurückentspannen. Für die Balmer-Reihe gilt: nfinal = 2.
In diesem Lab messen und beobachten Sie die Emissionsspektren von Wasserstoff, Helium und Neon und verwenden die Rydberg-Gleichung, um die Position der Spektrallinien zu bestimmen.

balmer serie

Das Bohr-Modell
Niels Bohr schlug 1913 ein Modell für das Wasserstoffatom vor, das beschrieb, dass diskrete Energiezustände mit einer festen Elektronenbahn um den Kern verbunden sind. Wichtig ist, dass ein Atom keine Energie entladen kann, solange sich seine Elektronen in stationären Zuständen befinden. Ein Elektron kann nur dann Energie abgeben, wenn es die Energiezustände ändert. Um Energiezustände zu ändern, muss sich ein Elektron von einer Umlaufbahn zur anderen bewegen, indem es entweder Energie absorbiert oder abgibt. Diese Änderung kann nur auftreten, wenn die aufgenommene oder abgegebene Energie der Differenz zwischen den beiden Zuständen entspricht. Elektronen können nicht zwischen den Umlaufbahnen existieren.
Die Quantenzahl n wird verwendet, um die verschiedenen Energiezustände zu markieren. Der niedrigste Energiezustand ist der Grundzustand, der n gleich eins ist. Die angeregten Zustände sind mit n gleich 2, 3, 4 usw. gekennzeichnet. Wenn ein Elektron im Grundzustand ein Photon absorbiert, dessen Energie der Differenz zwischen dem Grundzustand und dem zweiten Zustand entspricht, wird das Elektron angeregt und geht vom Grundzustand in den angeregten Zustand n= 2 über. Wenn die Energie des Photons gleich der Differenz zwischen dem Grundzustand und dem dritten Zustand ist, bewegt sich das Elektron in den Zustand n=3.
Nach dem Bohrschen Modell kann die potentielle Energie eines Elektrons in dern-ten Stufe mit der folgenden Gleichung berechnet werden:
<img src="/CDNSource/lm/labs/54/54_Concepts_1.jpg" alt="Bohr-Modell-Energieniveauformel, En=-Rhc/n², bezogen auf Elektronenübergänge, Gleichung." data-altupdatedat="1754907201000">
wobei En die potentielle Energie, R die Rydberg-Konstante (1,0974 × 107 m-1), h die Planck-Konstante (6,62607004 × 10-34 m2·kg/s) und c die Lichtgeschwindigkeit (~ 3 × 108 m/s) ist. Elektronen können auch spontan in den Grundzustand oder einen anderen niedriger angeregten Zustand zurückkehren. Wenn dies geschieht, wird die überschüssige Energie in Form eines emittierten Photons freigesetzt. Die Energie des Photons ist gleich der Energiedifferenz zwischen dem höheren und dem niedrigeren Energiezustand. Diese Energie entspricht den Wellenlängen des Lichts. Da jede Art von Atom ein unterschiedliches Energieniveau hat, variiert das Licht, das von jedem Übergang emittiert wird, für jedes Atom. Bei einer Probe gemischter Moleküle enthält das emittierte Licht einen Wellenlängenbereich in einem sogenannten kontinuierlichen Spektrum. Bei einer Probe, die Atome eines einzelnen Elements enthält, enthält das emittierte Licht nur bestimmte Wellenlängen, die als diskrete Linien betrachtet werden können, sobald sie durch ein Prisma getrennt sind.
<h4>Das Wasserstoffatom</h4>
Betrachtet man das Wasserstoffatom, so erfordert die Anregung seiner Elektronen die Absorption von ausreichend Energie, um die Bindung im zweiatomigen Molekül H2 aufzuspalten. Da für die Spaltung des Moleküls mehr Energie aufgewendet wird als benötigt, nehmen die Elektronen im Wasserstoffatom die überschüssige Energie auf und werden auf ein höheres Energieniveau angeregt. Wenn die Elektronen spontan auf ein niedrigeres Energieniveau zurückkehren, wird Licht emittiert, das der Energiedifferenz zwischen dem angeregten und dem niedrigeren Niveau entspricht.
Bei der Diskussion über die Emission von Energie wird das höhere Energieniveau als Anfangsniveau oder ni betrachtet, während das niedrigere Niveau als endgültiges Niveau oder nf betrachtet wird. Die Wellenlängen des emittierten Lichts hängen letztlich von der Energiedifferenz zwischen den beiden Niveaus ab.
In einer reinen Probe von Wasserstoffgas erscheint das Emissionsspektrum als deutliche Linien diskreter Wellenlängen, die spezifisch für das Element Wasserstoff sind. Einige dieser Linien liegen im sichtbaren Bereich des elektromagnetischen Spektrums, während andere im ultravioletten oder infraroten Bereich liegen.
<h4>Die Balmer-Serie</h4>
Die Reihe der sichtbaren Linien im Wasserstoffatomspektrum wird als Balmer-Reihe bezeichnet. Diese Reihe von spektralen Emissionslinien tritt auf, wenn das Elektron von einem hohen Energieniveau auf das niedrigere Energieniveau von n=2 übergeht. Johann Balmer beobachtete diese Spektrallinien bei 410,2 nm, 434,1 nm, 486,1 nm und 656,3 nm, was Übergängen von den Energieniveaus n=6, n=5, n=4 und n=3 zu den n=2 entspricht.
Balmer war in der Lage, diese Wellenlängen des emittierten Lichts mit Hilfe der Balmer-Formel in Beziehung zu setzen.
<img src="/CDNSource/lm/labs/54/54_Concepts_2.jpg" alt="Wellenlängenberechnungsformel λ=C(m²/(m²-n²)), Gleichung für die Spektroskopie-Analyse." data-altupdatedat="1754907201000">
Dabei ist λ die beobachtete Wellenlänge, C ist eine Konstante (364,50682 nm), n ist das niedrigere Energieniveau mit einem Wert von 2 und m ist das höhere Energieniveau, das einen Wert größer als 3 hat. Diese Beobachtung wurde dann von Johannes Rydberg verfeinert, wobei R die Rydberg-Konstante ist.
<img src="/CDNSource/lm/labs/54/54_Concepts_3.jpg" alt="Gleichung der Rydberg-Formel; berechnet Spektrallinien in der Atomspektroskopie." data-altupdatedat="1754907201000">
Denken Sie daran, dass diese Gleichung das emittierte Licht beschreibt, so dass das höhere Energieniveau als Anfangsniveau oder ni betrachtet wird, während das niedrigere Niveau als endgültiges Niveau oder nf betrachtet wird. Im Falle der Balmer-Reihe ist nf gleich 2. Diese Gleichung wurde mit dem Bohr-Modell kombiniert, um die Energie zu berechnen, die benötigt wird, um ein Elektron zwischen seinem Anfangs- und Endenergieniveau ΔE zu bewegen.
<img src="/CDNSource/lm/labs/54/54_Concepts_4.jpg" alt="Übergangsgleichung für das Energieniveau \(\Delta E = Rhc(\frac{1}{n_f^2} - \frac{1}{n_i^2})\); Physik-Formel." data-altupdatedat="1754907201000">
Später wurden weitere Spektralreihen für das Wasserstoffatom entdeckt. Die Lyman-Reihe enthält beispielsweise Emissionslinien mit Energien im ultravioletten Bereich.
<h4>Referenzen</h4>
<ol>
	<li style="font-size:17px">Kotz, J.C., Treichel Jr., P.M., Townsend, J.R. (2012). Chemie und chemische Reaktivität. Belmont, Kalifornien: Brooks/Cole, Cengage Lernen.</li>
	<li style="font-size:17px">Silderberg, M.S. (2009). Chemie: Die molekulare Natur von Materie und Veränderung. Boston, MA: McGraw Hill.</li>
</ol>

Balmer Serie

Das Bohr-Modell
Niels Bohr schlug 1913 ein Modell für das Wasserstoffatom vor, das beschrieb, dass diskrete Energiezustände mit einer festen Elektronenbahn verbunden sind ...

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hydrogen emission lines

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.

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.

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.

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.

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.

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.

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.

Hydrogen Emission Lines

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Earth’s Place in the Universe

HS-ESS1-2

south-carolina-college-and-career-ready-science-standards

South Carolina College and Career Ready Science Standards

Matter and Its Interactions

C-PS1-4

Georgia Standards of Excellence

9-12

SC1.

Light Emission and Electron Movement

Bereiten Sie den Unterricht zu Balmer Serie mit diesem 1-minütigen Leitfaden für Lehrkräfte vor. Behandelt die wichtigsten Vorbereitungsschritte für das Chemistry-Labor auf JoVE Lab Manual.

Balmer Serie Lehrervorbereitung in Chemistry - 2.0K+ Aufrufe | JoVE Lab Manual

Folgen Sie den Schritt-für-Schritt-Verfahren von Balmer Serie in diesem 4-minütigen Studentenprotokoll-Video. Teil der Chemistry-Sammlung auf JoVE Lab Manual.

Balmer Serie Studentenprotokoll in Chemistry - 5.1K+ Aufrufe | JoVE Lab Manual

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 ...

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 ...

Ions and Ionic Charges

In ordinary chemical reactions, the nucleus &mdash; which contains the protons and neutrons of each atom and thus identifies the element &mdash; remains ...

The Periodic Table

As early chemists discovered more elements, they realized that various elements could be grouped by their similar chemical behaviors. One such grouping ...

Atomic Mass

Atoms &mdash; and the protons, neutrons, and electrons that compose them &mdash; are extremely small. For example, a carbon atom weighs less than 2 ...

Molar Mass

The identity of a substance is defined not only by the types of atoms or ions it contains but by the quantity of each type of atom or ion. For example, ...

Periodic Classification of the Elements

The periodic table arranges atoms based on increasing atomic number so that elements with the same chemical properties recur periodically. When their ...

Types of Chemical Bonds

Chemical bonding theories were pioneered by American chemist Gilbert&#xA0;N. Lewis. He developed a model called the Lewis model to explain the type and ...

Ionic Bonding and Electron Transfer

Ions are atoms or molecules bearing an electrical charge. A cation (a positive ion) forms when a neutral atom loses one or more electrons from its valence ...

The Born-Haber Cycle

Lattice Energy&#xA0;

An ionic compound is stable because of the electrostatic attraction between its positive and negative ions. The lattice energy of a ...

trends-in-lattice-energy-ion-size-and-charge

Trends in Lattice Energy: Ion Size and Charge

An ionic compound is stable because of the electrostatic attraction between its positive and negative ions. The lattice energy of a compound is a measure ...

Covalent Bonding and Lewis Structures

Compared to ionic bonds, which results from the transfer of electrons between metallic and nonmetallic atoms, covalent bonds result from the mutual ...

Electronegativity

Whether a bond is nonpolar or polar covalent is determined by a property of the bonding atoms called electronegativity.&#xA0;

Electronegativity values of ...

bond-polarity-dipole-moment-and-percent-ionic-character

Bond Polarity, Dipole Moment, and Percent Ionic Character

Bond Polarity

The absolute value of the difference in electronegativity (&Delta;EN) of two bonded atoms provides a rough measure of the polarity to be ...

lewis-structures-of-molecular-compounds-and-polyatomic-ions

Lewis Structures of Molecular Compounds and Polyatomic Ions

To draw Lewis structures for complicated molecules and molecular ions, it is helpful to follow a step-by-step procedure as outlined:


	Determine the ...

Bond Energies and Bond Lengths

Stable molecules exist because covalent bonds hold the atoms together. The strength of a covalent bond is measured by the energy required to break it, ...

Bonding in Metals

Metallic bonds are formed between two metal atoms. A simplified model to describe metallic bonding has been developed by Paul Dr&#xFC;de called the ...

Model basic atomic structure and relate an element's atomic structure to its bonding, reactivity, and placement on the Periodic Table.

IPC.7.A

physical-and-chemical-properties-of-matter

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 ...

atomic-radii-and-effective-nuclear-charge

Atomic Radii and Effective Nuclear Charge

The elements in groups of the periodic table exhibit similar chemical behavior. This similarity occurs because the members of a group have the same number ...

Ionic Radii

Ionic radius is the measure used to describe the size of an ion. A cation always has fewer electrons and the same number of protons as the parent atom; it ...

Ionization Energy

The amount of energy required to remove the most loosely bound electron from a gaseous atom in its ground state is called its first ionization energy ...

Electron Affinity

The electron affinity (EA) is the energy change for adding an electron to a gaseous atom to form an anion (negative ion).



This process can be either ...

Alkali Metals

Group 1 elements are soft and shiny metallic solids. They are malleable, ductile, and good conductors of heat and electricity. The melting points of the ...

Use patterns within the Periodic Table to predict the relative physical and chemical properties of elements.

IPC.7.B

Science concepts. The student knows that relationships exist between the structure and properties of matter. The student is expected to:

Hydrogen Emission Lines

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