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

Los átomos y las moléculas exhiben un comportamiento muy interesante cuando absorben y liberan energía. Los electrones en un átomo existen tradicionalmente en su estado de energía más bajo, llamado estado fundamental, y etiquetado como n igual a uno. Sin embargo, cuando un átomo absorbe energía, los electrones se excitan y se mueven a un nivel de energía más alto. Cuando los electrones se relajan a un estado de energía más bajo o al estado fundamental, el exceso de energía se libera en forma de luz emitida.
Es posible que esté familiarizado con estos estados de energía, tal como se representan en el modelo de Bohr, que describe un átomo como un núcleo con electrones en órbita en capas u orbitales. Estas capas son las mismas que los niveles de energía y están etiquetadas por n. La longitud de onda de la luz emitida depende de la diferencia entre los niveles de energía alta y baja. La luz emitida de alta energía es el resultado de la relajación de los electrones desde un nivel de energía más alto, y la luz emitida de baja energía es el resultado de la relajación de los electrones desde un nivel de energía más bajo.
El espectro de emisión es una medida de la radiación emitida a través de un rango de longitudes de onda. Con las especies elementales puras, el comportamiento de emisión aparece como líneas de longitudes de onda específicas en lugar de un espectro amplio. Dado que los diferentes átomos tienen diferentes niveles de energía, estas líneas espectrales varían de un elemento a otro y dependen de las transiciones que esos electrones hacen entre los estados de energía cuando se excitan. Por ejemplo, hay seis series de líneas espectrales con nombre para el hidrógeno, una de las cuales es la serie de Balmer.
La serie de Balmer de líneas espectrales ocurre cuando los electrones pasan de un nivel de energía superior a n = 3 a n = 2. El espectro de luz visible para la serie Balmer aparece como líneas espectrales a 410, 434, 486 y 656 nm. La línea h alfa es la línea roja a 656 nm y se produce debido a la transición de n= 3 a n= 2. Las líneas cian, azul y violeta corresponden a las transiciones de nivel de energía de n = 4, 5 y 6, respectivamente, hasta n = 2. Se pueden medir líneas espectrales adicionales fuera del rango visible.
Johann Balmer cuantificó las líneas visibles utilizando la fórmula de Balmer. Aquí, lambda es la longitud de onda observada, c es una constante, n es el nivel de energía más bajo de dos y m es el nivel de energía más alto. La combinación de la ecuación de Balmer y el modelo de Bohr nos da la ecuación de Rydberg, que describe las líneas espectrales de muchos elementos diferentes. En esta ecuación, lambda es la longitud de onda registrada y RH es la constante de Rydberg. Los niveles iniciales marcados con n-inicial representan el nivel de energía más alto al que se excitan los electrones y n final es el nivel de energía más bajo al que los electrones se relajan. Para la serie Balmer, nfinal = 2.
En este laboratorio, medirá y observará los espectros de emisión de hidrógeno, helio y neón y utilizará la ecuación de Rydberg para determinar la ubicación de las líneas espectrales.

serie balmer

El modelo de Bohr
Niels Bohr propuso un modelo para el átomo de hidrógeno en 1913 que describía que los estados de energía discretos están asociados con una órbita fija de electrones alrededor del núcleo. Es importante destacar que un átomo no puede descargar energía mientras sus electrones están en estado estacionario. Un electrón solo puede emitir energía cambiando los estados de energía. Para cambiar los estados de energía, un electrón debe moverse de una órbita a otra absorbiendo o emitiendo energía. Este cambio solo puede ocurrir si la energía absorbida o emitida es igual a la diferencia entre los dos estados. Los electrones no pueden existir entre órbitas.
El número cuántico, n, se utiliza para etiquetar los diferentes estados de energía. El estado de energía más bajo es el estado fundamental, que es n igual a uno. Los estados excitados se etiquetan como n igual a 2, 3, 4, y así sucesivamente. Cuando un electrón en el estado fundamental absorbe un fotón cuya energía es igual a la diferencia entre el estado fundamental y el segundo estado, el electrón se excita y pasa del estado fundamental al estado excitado n= 2. Si la energía del fotón es igual a la diferencia entre el estado fundamental y el tercero, el electrón pasa al estado n=3.
De acuerdo con el modelo de Bohr, la energía potencial de un electrón en el nivel n se puede calcular utilizando la siguiente ecuación:
<img src="/CDNSource/lm/labs/54/54_Concepts_1.jpg" alt="Fórmula del nivel de energía del modelo de Bohr, En=-Rhc/n², relacionada con las transiciones de electrones, ecuación." data-altupdatedat="1754907201000">
donde En es la energía potencial, R es la constante de Rydberg (1,0974 × 107 m-1), h es la constante de Planck (6,62607004 × 10-34 m2 kg/s) y c es la velocidad de la luz (~ 3 × 108 m/s). Los electrones también pueden regresar espontáneamente al estado fundamental o a cualquier otro estado excitado inferior. Cuando esto sucede, el exceso de energía se libera en forma de un fotón emitido. La energía del fotón es igual a la diferencia de energía entre los estados de energía más altos y más bajos. Esa energía corresponde a las longitudes de onda de la luz. Dado que cada tipo de átomo tiene diferentes niveles de energía, la luz emitida por cada transición varía para cada átomo. Para una muestra de moléculas mixtas, la luz emitida contiene un rango de longitudes de onda en lo que se llama un espectro continuo. Para una muestra que contiene átomos de un solo elemento, la luz emitida contiene solo ciertas longitudes de onda, que se pueden ver como líneas discretas una vez separadas por un prisma.
<h4>El átomo de hidrógeno</h4>
Mirando específicamente al átomo de hidrógeno, la excitación de sus electrones requiere la absorción de suficiente energía para dividir el enlace en la molécula diatómicaH2. Dado que se utiliza más energía de la necesaria para dividir la molécula, los electrones del átomo de hidrógeno absorben el exceso de energía y se excitan a un nivel de energía más alto. Cuando los electrones regresan espontáneamente a un nivel de energía más bajo, se emite luz, que corresponde a la diferencia de energía entre el nivel excitado y el nivel inferior.
Cuando se habla de la emisión de energía, el nivel de energía más alto se considera el nivel inicial, o ni, mientras que el nivel inferior se considera el nivel final, o nf. Las longitudes de onda de la luz emitida dependen en última instancia de la diferencia de energía entre los dos niveles.
En una muestra pura de gas hidrógeno, el espectro de emisión aparece como líneas distintas de longitudes de onda discretas que son específicas del elemento hidrógeno. Algunas de estas líneas se encuentran en el rango visible del espectro electromagnético, mientras que otras están en el rango ultravioleta o infrarrojo.
<h4>La serie Balmer</h4>
La serie de líneas visibles en el espectro del átomo de hidrógeno se denomina serie de Balmer. Esta serie de líneas de emisión espectral ocurre cuando el electrón pasa de un nivel de alta energía al nivel de energía más bajo de n = 2. Johann Balmer observó estas líneas espectrales a 410,2 nm, 434,1 nm, 486,1 nm y 656,3 nm, que corresponden a transiciones de los niveles de energía n=6, n=5, n=4 y n=3 al nivel n=2, respectivamente.
Balmer fue capaz de relacionar estas longitudes de onda de la luz emitida utilizando la fórmula de Balmer.
<img src="/CDNSource/lm/labs/54/54_Concepts_2.jpg" alt="Fórmula de cálculo de longitud de onda λ=C(m²/(m²-n²)), ecuación para análisis espectroscópico." data-altupdatedat="1754907201000">
Aquí, λ es la longitud de onda observada, C es una constante (364,50682 nm), n es el nivel de energía más bajo con un valor de 2 y m es el nivel de energía más alto, que tiene un valor mayor que 3. Esta observación fue refinada por Johannes Rydberg, donde R es la constante de Rydberg.
<img src="/CDNSource/lm/labs/54/54_Concepts_3.jpg" alt="ecuación de la fórmula de Rydberg; calcula líneas espectrales en espectroscopia atómica." data-altupdatedat="1754907201000">
Recuerde que esta ecuación describe la luz emitida, por lo que el nivel de energía más alto se considera el nivel inicial, o ni, mientras que el nivel inferior se considera el nivel final, o nf. En el caso de la serie de Balmer, nf es igual a 2. Esta ecuación se combinó con el modelo de Bohr para calcular la energía necesaria para mover un electrón entre sus niveles de energía inicial y final, ΔE.
<img src="/CDNSource/lm/labs/54/54_Concepts_4.jpg" alt="Ecuación de transición de nivel de energía \(\Delta E = Rhc(\frac{1}{n_f^2} - \frac{1}{n_i^2})\); Fórmula física." data-altupdatedat="1754907201000">
Más tarde, se descubrieron otras series espectrales para el átomo de hidrógeno. Por ejemplo, la serie Lyman contiene líneas de emisión con energías en la región ultravioleta.
<h4>Referencias</h4>
<ol>
	<li style="font-size:17px">Kotz, J.C., Treichel Jr, P.M., Townsend, J.R. (2012). Química y reactividad química. Belmont, CA: Brooks/Cole, Cengage Learning.</li>
	<li style="font-size:17px">Silderberg, M.S. (2009). Química: La naturaleza molecular de la materia y el cambio. Boston, MA: McGraw Hill.</li>
</ol>

Serie Balmer

El modelo de Bohr
Niels Bohr propuso un modelo para el átomo de hidrógeno en 1913 que describía que los estados de energía discretos están asociados con una órbita fija de electrones.

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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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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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