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

Quando un atomo assorbe energia, gli elettroni si eccitano e si spostano a un livello di energia superiore. Quando gli elettroni si rilassano fino a uno stato di energia inferiore o allo stato fondamentale, l'energia in eccesso viene rilasciata come un fotone. La lunghezza d'onda della luce assorbita ed emessa dipende dalla differenza tra i livelli di energia alta e bassa.

La luce emessa ad alta energia è il risultato di elettroni che si rilassano da un livello di energia più elevato, e la luce emessa a bassa energia è il risultato di elettroni che si rilassano da un livello di energia più basso. Uno 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.

Questo è lo spettro di emissione dell'idrogeno. L'insieme di linee spettrali nella regione della luce visibile è noto come serie di Balmer. Esso si verifica quando gli elettroni passano da un livello di energia superiore a n 3 fino a n 2.

Lo spettro della luce visibile appare come linee spettrali a 410, 434, 486 e 656 nm, che corrispondono alle transizioni del livello di energia da n 3, 4, 5 e 6, rispettivamente, a n 2. È possibile misurare ulteriori linee spettrali al di fuori del campo visibile, come la serie Lyman nella regione UV e la serie Paschen nella regione infrarossa. Le lunghezze d'onda delle linee spettrali per l'idrogeno possono essere previste utilizzando un'espressione matematica, dove R-H è la costante di Rydberg, n1 è il numero quantico principale del livello di energia inferiore e n2 è il numero quantico principale per il livello di energia superiore.

Per la serie di Balmer, n1 2. Poiché atomi diversi hanno livelli di energia diversi, le linee di emissione spettrale variano da elemento a elemento e vengono utilizzate per identificare le sostanze. L'inverso di uno spettro di emissione è il suo spettro di assorbimento.

Guardando l'idrogeno, le linee nel suo spettro di assorbimento si trovano alle stesse lunghezze d'onda del suo spettro di emissione, ma sono scure. Queste sono le lunghezze d'onda della luce che vengono assorbite da un atomo di idrogeno quando viene esposto a uno spettro di luce bianca continua.

spettri di emissione

Quando i solidi, i liquidi o i gas condensati vengono riscaldati a sufficienza, irradiano parte dell’energia in eccesso sotto forma di luce. I fotoni prodotti in questo modo hanno una gamma di energie e quindi producono uno spettro continuo in cui è presente una serie ininterrotta di lunghezze d'onda.

A differenza degli spettri continui, la luce può anche presentarsi come spettri discreti o a linee con larghezze di linea molto strette intervallate da tutte le regioni spettrali. Eccitando un gas a bassa pressione parziale utilizzando una corrente elettrica, o riscaldandolo, si produrranno spettri a righe. Le lampadine fluorescenti e le insegne al neon funzionano in questo modo. Ogni elemento mostra il proprio caratteristico insieme di linee, così come le molecole, sebbene i loro spettri siano generalmente molto più complicati.

Ciascuna riga di emissione è costituita da un'unica lunghezza d'onda della luce, il che implica che la luce emessa da un gas è costituita da un insieme di energie discrete. Ad esempio, quando una scarica elettrica passa attraverso un tubo contenente idrogeno gassoso a bassa pressione, le molecole di H_2 vengono scomposte in atomi di H separati e si osserva un colore blu-rosa. Il passaggio della luce attraverso un prisma produce uno spettro a linee, indicando che questa luce è composta da fotoni di quattro lunghezze d'onda visibili.

L'origine degli spettri discreti negli atomi e nelle molecole era estremamente sconcertante per gli scienziati della fine del XIX secolo. Secondo la teoria elettromagnetica classica, dovrebbero essere osservati solo spettri continui. Altre linee discrete per l'atomo di idrogeno sono state trovate nelle regioni UV e IR. Johannes Rydberg generalizzò il lavoro di Balmer e sviluppò una formula empirica che prevedeva tutte le linee di emissione dell'idrogeno, non solo quelle limitate all'intervallo visibile, dove n_1 e n_2 sono numeri interi, n_1 < n_2

<img alt="Eq1" src="https://www.jove.com/files/core/CDNSource/lessons/977/977_Equations-01.png" style="height: 40px;" />

Anche alla fine del XIX secolo, la spettroscopia era una scienza molto precisa, e quindi le lunghezze d'onda dell'idrogeno venivano misurate con altissima precisione, il che implicava che anche la costante di Rydberg potesse essere determinata in modo molto preciso. Che una formula così semplice come quella di Rydberg potesse spiegare misurazioni così precise sembrava sorprendente all'epoca, ma fu la spiegazione finale degli spettri di emissione data da Neils Bohr nel 1913 che alla fine convinse gli scienziati ad abbandonare la fisica classica e stimolò lo sviluppo della moderna formula quantistica. meccanica.

Questo testo è adattato da <a href="https://openstax.org/books/chemistry-atoms-first-2e/pages/3-1-electromagnetic-energy">Openstax, Chemistry 2e, Section 3.1: Electromagnetic Energy</a>.

Spettri di emissione

Quando i solidi, i liquidi o i gas condensati vengono riscaldati a sufficienza, irradiano parte dell'energia in eccesso sotto forma di luce. I fotoni prodotti in questo modo hanno un ...

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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

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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...

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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...

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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...

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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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Line Spectra from Heated Gases

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