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

K12 Standards

Texas Essential Knowledge and Skills

High School

Chemistry and Physics

IPC.7

IPC.7.E

Balmer Series

Gli atomi e le molecole mostrano un comportamento molto interessante quando assorbono e rilasciano energia. Gli elettroni in un atomo esistono tradizionalmente al loro stato energetico più basso, chiamato stato fondamentale, ed etichettati come n uguale a uno. Tuttavia, quando un atomo assorbe energia, gli elettroni si eccitano e si spostano a un livello di energia più elevato. Quando gli elettroni si rilassano fino a uno stato di energia inferiore o allo stato fondamentale, l'energia in eccesso viene rilasciata sotto forma di luce emessa.
Potresti avere familiarità con questi stati energetici, poiché sono raffigurati nel modello di Bohr, che descrive un atomo come un nucleo con elettroni orbitanti in gusci o orbitali. Questi gusci sono gli stessi dei livelli di energia e sono etichettati con n. La lunghezza d'onda della luce emessa dipende dalla differenza tra i livelli di energia alta e bassa. La luce ad alta energia emessa deriva dal rilassamento degli elettroni da un livello di energia più elevato e la luce a bassa emissione di energia deriva dal rilassamento degli elettroni da un livello di energia inferiore.
Lo spettro di emissione è una misura della radiazione emessa in una gamma di lunghezze d'onda. Con le specie elementari pure, il comportamento di emissione appare come linee di lunghezze d'onda specifiche piuttosto che come un ampio spettro. Poiché atomi diversi hanno livelli di energia diversi, queste linee spettrali variano da elemento a elemento e dipendono dalle transizioni che quegli elettroni compiono tra gli stati di energia quando sono eccitati. Ad esempio, ci sono sei serie di linee spettrali per l'idrogeno, una delle quali è la serie di Balmer.
La serie di linee spettrali di Balmer si verifica quando gli elettroni passano da un livello di energia superiore a n = 3 fino a n = 2. Lo spettro della luce visibile per la serie di Balmer appare come linee spettrali a 410, 434, 486 e 656 nm. La riga h alfa è la linea rossa a 656 nm e si verifica a causa della transizione da n= 3 a n = 2. Le linee ciano, blu e viola corrispondono alle transizioni del livello di energia da n = 4, 5 e 6, rispettivamente, fino a n = 2. Ulteriori linee spettrali possono essere misurate al di fuori dell'intervallo visibile.
Johann Balmer ha quantificato le linee visibili utilizzando la formula di Balmer. Qui, lambda è la lunghezza d'onda osservata, c è una costante, n è il livello di energia inferiore di due e m è il livello di energia più alto. La combinazione dell'equazione di Balmer e del modello di Bohr ci dà l'equazione di Rydberg, che descrive le linee spettrali di molti elementi diversi. In questa equazione, lambda è la lunghezza d'onda registrata e RH è la costante di Rydberg. I livelli iniziali contrassegnati da n-iniziale rappresentano il livello di energia più alto a cui gli elettroni eccitano e n finale è il livello di energia più basso a cui gli elettroni si rilassano. Per la serie di Balmer, nfinale = 2.
In questo laboratorio, misurerai e osserverai gli spettri di emissione di idrogeno, elio e neon e utilizzerai l'equazione di Rydberg per determinare la posizione delle linee spettrali.

serie balmer

Il modello Bohr
Niels Bohr propose un modello per l'atomo di idrogeno nel 1913 che descriveva gli stati di energia discreti associati a un'orbita elettronica fissa attorno al nucleo. È importante sottolineare che un atomo non può scaricare energia mentre i suoi elettroni sono in stato stazionario. Un elettrone può emettere energia solo cambiando gli stati energetici. Per cambiare gli stati energetici, un elettrone deve spostarsi da un'orbita all'altra assorbendo o emettendo energia. Questo cambiamento può avvenire solo se l'energia assorbita o emessa è uguale alla differenza tra i due stati. Gli elettroni non possono esistere tra le orbite.
Il numero quantico, n, viene utilizzato per etichettare i diversi stati energetici. Lo stato energetico più basso è lo stato fondamentale, che è n uguale a uno. Gli stati eccitati sono etichettati n uguale a 2, 3, 4 e così via. Quando un elettrone allo stato fondamentale assorbe un fotone la cui energia è uguale alla differenza tra lo stato fondamentale e il secondo stato, l'elettrone si eccita e passa dallo stato fondamentale allo stato eccitato n= 2. Se l'energia del fotone è uguale alla differenza tra lo stato fondamentale e il terzo, l'elettrone si sposta allo stato n=3.
Secondo il modello di Bohr, l'energia potenziale di un elettrone nel livello n° può essere calcolata utilizzando la seguente equazione:
<img src="/CDNSource/lm/labs/54/54_Concepts_1.jpg" alt="Formula del livello di energia del modello di Bohr, En=-Rhc/n², relativa alle transizioni elettroniche, equazione." data-altupdatedat="1754907201000">
dove En è l'energia potenziale, R è la costante di Rydberg (1,0974 × 107 m-1), h è la costante di Planck (6,62607004 × 10-34 m2·kg/s) e c è la velocità della luce (~ 3 × 108 m/s). Gli elettroni possono anche tornare spontaneamente allo stato fondamentale o a qualsiasi altro stato eccitato inferiore. Quando ciò accade, l'energia in eccesso viene rilasciata sotto forma di un fotone emesso. L'energia del fotone è uguale alla differenza di energia tra gli stati di energia superiore e inferiore. Quell'energia corrisponde alle lunghezze d'onda della luce. Poiché ogni tipo di atomo ha livelli di energia diversi, la luce emessa da ogni transizione varia per ogni atomo. Per un campione di molecole miste, la luce emessa contiene una gamma di lunghezze d'onda in quello che viene chiamato spettro continuo. Per un campione contenente atomi di un singolo elemento, la luce emessa contiene solo determinate lunghezze d'onda, che possono essere visualizzate come linee discrete una volta separate da un prisma.
<h4>L'atomo di idrogeno</h4>
Guardando in particolare all'atomo di idrogeno, l'eccitazione dei suoi elettroni richiede l'assorbimento di energia sufficiente a dividere il legame nella molecola biatomica H2. Poiché viene utilizzata più energia del necessario per scindere la molecola, gli elettroni nell'atomo di idrogeno assorbono l'energia in eccesso e vengono eccitati a un livello di energia più elevato. Quando gli elettroni ritornano spontaneamente a un livello di energia inferiore, viene emessa luce, che corrisponde alla differenza di energia tra il livello eccitato e il livello inferiore.
Quando si discute dell'emissione di energia, il livello di energia più alto è considerato il livello iniziale, o ni, mentre il livello più basso è considerato il livello finale, o nf. Le lunghezze d'onda della luce emessa dipendono in ultima analisi dalla differenza di energia tra i due livelli.
In un campione puro di idrogeno gassoso, lo spettro di emissione appare come linee distinte di lunghezze d'onda discrete specifiche per l'elemento idrogeno. Alcune di queste linee si trovano nella gamma visibile dello spettro elettromagnetico, mentre altre si trovano nella gamma dell'ultravioletto o dell'infrarosso.
<h4>La serie Balmer</h4>
La serie di linee visibili nello spettro dell'atomo di idrogeno è chiamata serie di Balmer. Questa serie di righe di emissione spettrale si verifica quando l'elettrone passa da un livello di energia elevato a un livello di energia inferiore di n=2. Johann Balmer ha osservato queste linee spettrali a 410,2 nm, 434,1 nm, 486,1 nm e 656,3 nm, che corrispondono alle transizioni dai livelli di energia n=6, n=5, n=4 e n=3 al livello n=2, rispettivamente.
Balmer è stato in grado di mettere in relazione queste lunghezze d'onda della luce emessa utilizzando la formula di Balmer.
<img src="/CDNSource/lm/labs/54/54_Concepts_2.jpg" alt="Formula di calcolo della lunghezza d'onda λ=C(m²/(m²-n²)), equazione per l'analisi spettroscopica." data-altupdatedat="1754907201000">
Qui, λ è la lunghezza d'onda osservata, C è una costante (364,50682 nm), n è il livello di energia inferiore con un valore di 2 e m è il livello di energia più alto, che ha un valore maggiore di 3. Questa osservazione è stata poi perfezionata da Johannes Rydberg, dove R è la costante di Rydberg.
<img src="/CDNSource/lm/labs/54/54_Concepts_3.jpg" alt="equazione della formula di Rydberg; Calcola le linee spettrali nella spettroscopia atomica." data-altupdatedat="1754907201000">
Ricorda che questa equazione descrive la luce emessa, quindi il livello di energia più alto è considerato il livello iniziale, o ni, mentre il livello più basso è considerato il livello finale, o nf. Nel caso della serie di Balmer, nf è uguale a 2. Questa equazione è stata combinata con il modello di Bohr per calcolare l'energia necessaria per spostare un elettrone tra i suoi livelli di energia iniziale e finale, ΔE.
<img src="/CDNSource/lm/labs/54/54_Concepts_4.jpg" alt="Equazione di transizione del livello di energia \(\Delta E = Rhc(\frac{1}{n_f^2} - \frac{1}{n_i^2})\); formula di fisica." data-altupdatedat="1754907201000">
Più tardi, sono state scoperte altre serie spettrali per l'atomo di idrogeno. Ad esempio, la serie di Lyman contiene linee di emissione con energie nella regione dell'ultravioletto.
<h4>referenze</h4>
<ol>
	<li style="font-size:17px">Kotz, J.C., Treichel Jr, P.M., Townsend, J.R. (2012). Chimica e reattività chimica. Belmont, CA: Brooks/Cole, Cengage Learning.</li>
	<li style="font-size:17px">Silderberg, M.S. (2009). Chimica: la natura molecolare della materia e il cambiamento. Boston, MA: McGraw Hill.</li>
</ol>

Serie Balmer

Il modello Bohr
Niels Bohr propose un modello per l'atomo di idrogeno nel 1913 che descriveva gli stati di energia discreti associati a un'orbita di elettroni fissa ...

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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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Explore how hydrogen electrons jump between energy levels and produce discrete emission lines, including the visible Balmer series and other hydrogen spectra.

Hydrogen Emission Lines | JoVE K12

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Preparati a insegnare Serie Balmer con questa guida per istruttori di 1 minuti. Copre i passaggi chiave di preparazione per il laboratorio Chemistry su JoVE Lab Manual.

Preparazione dell'istruttore per Serie Balmer in Chemistry - Oltre 2.0K visualizzazioni | JoVE Lab Manual

Segui le procedure passo dopo passo di Serie Balmer in questo video di protocollo studente di 4 minuti. Parte della collezione Chemistry su JoVE Lab Manual.

Protocollo studente per Serie Balmer in Chemistry - Oltre 5.1K visualizzazioni | 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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