Name: shielding and orbital penetration
Uploaded: 2020-09-03T13:38:28.000Z
Duration: 3 min 21 s
Channel: JoVE

K12 Standards

Texas Essential Knowledge and Skills

High School

Chemistry and Physics

IPC.7

IPC.7.E

The Energies of Atomic Orbitals

Los orbitales atómicos tienen diferentes energías, racionalizadas por las interacciones de Coulomb, el efecto de blindaje y la penetración orbital. La ley de Coulomb indica que la fuerza atractiva o repulsiva entre dos partículas cargadas tiene una relación de cuadrado inverso con la distancia entre ellos. Los tamaños de los orbitales atómicos aumentan con el número de proyectiles, y los electrones se repelen del espacio ocupado por los orbitales de la capa inferior.Por lo tanto, la ley de Coulomb sugiere que como el número de la capa aumenta, los electrones experimentan menos atracción al núcleo, que corresponde a las energías orbitales superiores. Además, los electrones que están aproximadamente a la misma distancia desde o más cerca del núcleo tienen un efecto protector que reduce aún más la atracción hacia el núcleo. Cuanto mayor es el blindaje, menor es la atracción hacia el núcleo.Esta es una de las razones de las diferencias en energías orbitales dentro de las capas de electrones. Por ejemplo, los electrones 3s y 3p protegen significativamente a los electrones 3d. La carga nuclear efectiva que siente un electrón se calcula restando la constante de apantallamiento S, que depende de la cantidad de electrones de protección y las subcapas que ocupan, a partir del número atómico.Por ejemplo, los dos electrones 1s en el litio, que tiene un número atómico de tres, oculta su electrón 2s. La constante de blindaje para ese electrón se determina a partir de reglas semi-empíricas y la misma es 1.7. Por lo tanto, la carga nuclear efectiva que siente el electrón 2s es 1, 3.Las formas de los orbitales también dictan su energía. Si los electrones en un orbital externo pueden moverse lejos en áreas ocupadas por electrones internos para estar cerca del núcleo, ellos estarán mucho menos protegidos allí. Por lo tanto, la energía de ese orbital externo es menor.Esto se puede visualizar con una función de distribución radial. que describe la probabilidad de encontrar un electrón a una distancia del núcleo determinada. Los gráficos de la función de distribución radial para las subcapas 1, 2 y 2p revelan que los electrones 2s tienen una probabilidad modesta de estar cerca del núcleo, mientras que los electrones 2p permanecen principalmente en el borde exterior de la región 1s.El orbital 2s, por lo tanto, se dice que tiene mayor capacidad de penetración. En la tercera capa, los electrones 3s penetran más y los electrones 3d penetran menos. Generalmente, la energía orbital de átomo aumenta con el número de capas.y, a nivel de subcapa, de s a f. Sin embargo, el efecto de penetración se vuelve tan significativo en la cuarta y quinta capa que los orbitales 4s y 5s con frecuencia tienen energías relativas más bajas que los orbitales 3d y 4d, respectivamente.

las energías de los orbitales atómicos

En un átomo, los electrones cargados negativamente son atraídos al núcleo cargado positivamente. En un átomo multielectrón, también se observan repulsiones electrón-electrón. Las fuerzas de atracción y repulsión dependen de la distancia entre las partículas, así como del signo y la magnitud de las cargas sobre las partículas individuales. Cuando las cargas sobre las partículas son opuestas, se atraen entre sí. Si ambas partículas tienen la misma carga, se repelen entre sí.
 
 A medida que aumenta la magnitud de las cargas, aumenta la magnitud de la fuerza. Sin embargo, cuando la separación de cargas es mayor, las fuerzas disminuyen. Así, la fuerza de atracción entre un electrón y su núcleo es directamente proporcional a la distancia entre ellos. Si el electrón está más cerca del núcleo, está más unido al núcleo; por lo tanto, los electrones en las diferentes capas (a diferentes distancias) tienen energías diferentes.
 
 Para los átomos con múltiples niveles de energía, los electrones internos protegen parcialmente a los electrones externos de la tracción del núcleo, debido a las repulsiones electrón-electrón. Los electrones centrales protegen a los electrones en las capas externas, mientras que los electrones en la misma capa de valencia no bloquean la atracción nuclear experimentada entre ellos de manera tan eficiente. Esto se puede explicar con el concepto de carga nuclear efectiva, Zef. Esta es la tracción ejercida por el núcleo sobre un electrón específico, teniendo en cuenta cualquier repulsión electrón-electrón. Para el hidrógeno, sólo hay un electrón, por lo que la carga nuclear (Z) y la carga nuclear efectiva (Zef) son iguales. Para todos los demás átomos, los electrones internos protegen parcialmente a los electrones externos de la tracción del núcleo, y así:
 
 <img alt="Eq1" src="https://www.jovecore.com/CDNSource/lessons/985/985_Equations-01.png" style="height: 40px;" />
 
 La penetración orbital describe la capacidad de un electrón para estar más cerca del núcleo. Los electrones en el orbital s pueden acercarse al núcleo y tener una capacidad más penetrante. La densidad de probabilidad para un orbital s esférico es distinta de cero en el núcleo.  Las distintas subcapas tienen orientaciones espaciales diferentes. Debido al orbital en forma de mancuerno, el electrón p penetra mucho menos. Su función de onda tiene un nodo que pasa a través del núcleo, donde la probabilidad de encontrar el electrón es cero. Por lo tanto, un electrón en el orbital s está unido más estrechamente al núcleo y tiene menos energía que el electrón p. Un electrón d tiene una penetración aún menor y una energía mayor que un electrón en el orbital p. 
 Para varias capas y subcapas, la tendencia del poder de penetración de un electrón se puede representar de la siguiente manera
 
 <img alt="Eq2" src="https://www.jovecore.com/CDNSource/lessons/985/985_Equations-02.png" style="height: 40px;" />
 
 El efecto del blindaje y la penetración es grande, y un electrón 4s puede tener una energía más baja que un electrón 3d.
 
 Este texto es adaptado de <a href="https://openstax.org/books/chemistry-2e/pages/6-4-electronic-structure-of-atoms-electron-configurations">Openstax, Química 2e, Sección 6.4: Estructura Electrónica de los Átomos (Configuraciones Electrónicas).</a>

Las Energías de Los Orbitales Atómicos

En un átomo, los electrones cargados negativamente son atraídos por el núcleo cargado positivamente. En un átomo multielectrón, las repulsiones electrón-electrón son ...

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KEY TERMS AND CONCEPTS

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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 electrons in different shells can have different energies because they sit at different distances from the nucleus. In atoms with several energy levels, inner electrons partly block the nucleus from outer electrons. This blocking is called shielding, and it happens because of electron-electron repulsions.

The pull felt by one electron after shielding is the effective nuclear charge, or Zeff. It describes the attraction from the nucleus after repulsions are taken into account. For hydrogen, which has only one electron, Z and Zeff are the same. In atoms with more than one electron, core electrons shield outer electrons, while electrons in the same valence shell do not block the nuclear pull very well.

Orbital penetration describes how close an electron can get to the nucleus. s orbitals penetrate best because their spherical probability density is non-zero at the nucleus. p orbitals penetrate less because their dumbbell shape has a node through the nucleus, where the chance of finding an electron is zero. d orbitals penetrate even less than p orbitals.

Because of shielding and penetration, s electrons are usually bound more tightly and have lower energy than p electrons. d electrons have higher energy than p electrons. These effects can be strong enough that a 4s electron may have lower energy than a 3d electron.

Shielding and Orbital Penetration

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the energies of atomic orbitals

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

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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

Learn how shielding, effective nuclear charge, and orbital penetration affect electron energy levels, and why s, p, and d orbitals differ in energy.

Shielding and Orbital Penetration | JoVE K12

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PS.Chm.1.3

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IPC.7.C

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Shielding and Orbital Penetration

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