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

Gli orbitali atomici hanno energie diverse, come razionalizzato dalle interazioni di Coulomb, dall'effetto schermante e dalla penetrazione orbitale. La legge di Coulomb indica che la forza attrattiva o repulsiva fra due particelle cariche ha una relazione pari all'inverso del quadrato della distanza tra loro. Le dimensioni degli orbitali atomici aumentano con il numero dei gusci e gli elettroni vengono respinti dallo spazio occupato dagli orbitali dei gusci inferiori.

Pertanto, la legge di Coulomb suggerisce che all'aumentare del numero di gusci, gli elettroni sperimentino meno attrazione per il nucleo, che corrisponde a energie orbitali più elevate. Inoltre, gli elettroni che si trovano all'incirca alla stessa distanza o più vicini al nucleo hanno un effetto di schermatura che riduce ulteriormente l'attrazione per il nucleo. Maggiore è la schermatura, minore è l'attrazione per il nucleo.

Questa è una delle spiegazioni per le differenze nelle energie orbitali all'interno dei gusci degli elettroni. Per esempio, gli elettroni 3s e 3p schermano in modo significativo gli elettroni 3d. La carica nucleare effettiva percepita da un elettrone si calcola sottraendo dal numero atomico la costante di schermatura S, che dipende dal numero di elettroni di schermatura e dai subshell che occupano.

Per esempio, i due elettroni 1s nel litio, avente un numero atomico di tre, schermano il suo elettrone 2s. La costante di schermatura per quell'elettrone è determinata da regole semi-empiriche pari a 1, 7. Pertanto, la carica nucleare effettiva percepita dall'elettrone 2s è 1.3.

Anche le forme degli orbitali determinano la loro energia. Se gli elettroni in un orbitale esterno possono spostarsi in aree occupate da elettroni interni per essere vicini al nucleo, saranno molto meno schermati. Pertanto, l'energia di quell'orbitale esterno è inferiore.

Questo può essere visualizzato con una funzione di distribuzione radiale che descrive la probabilità di trovare un elettrone a una data distanza dal nucleo. I grafici della funzione di distribuzione radiale per i subshell 1s, 2s e 2p rivelano che gli elettroni 2s hanno una probabilità modesta di essere vicini al nucleo, mentre gli elettroni 2p rimangono per lo più all'esterno o sul bordo esterno della regione 1s. L'orbitale 2s, dunque, si dice avere una maggiore capacità di penetrazione.

Nel terzo guscio, gli elettroni 3s penetrano di più e gli elettroni 3d penetrano di meno. Generalmente, l'energia orbitale atomica aumenta con il numero di gusci e, a livello di subshell, da s a f. Tuttavia, l'effetto di penetrazione diventa così significativo nel quarto e quinto guscio, che gli orbitali 4s e 5s hanno frequentemente energie relative inferiori, rispettivamente, degli orbitali 3d e 4d.

le energie degli orbitali atomici

In un atomo gli elettroni carichi negativamente sono attratti dal nucleo carico positivamente. In un atomo multielettronico si osservano anche repulsioni elettrone-elettrone. Le forze attrattive e repulsive dipendono dalla distanza tra le particelle, nonché dal segno e dall'entità delle cariche sulle singole particelle. Quando le cariche delle particelle sono opposte, si attraggono. Se entrambe le particelle hanno la stessa carica si respingono.

All’aumentare dell’entità delle cariche, aumenta l’entità della forza. Tuttavia, quando la separazione delle cariche è maggiore, le forze diminuiscono. Pertanto, la forza di attrazione tra un elettrone e il suo nucleo è direttamente proporzionale alla distanza tra loro. Se l'elettrone è più vicino al nucleo, è legato più strettamente al nucleo; pertanto, gli elettroni nei diversi gusci (a distanze diverse) hanno energie diverse.

Per gli atomi con più livelli energetici, gli elettroni interni schermano parzialmente gli elettroni esterni dall'attrazione del nucleo, a causa delle repulsioni elettrone-elettrone. Gli elettroni del nucleo proteggono gli elettroni nei gusci esterni, mentre gli elettroni nello stesso guscio di valenza non bloccano l'attrazione nucleare reciprocamente sperimentata in modo altrettanto efficiente. Ciò può essere spiegato con il concetto di carica nucleare effettiva, Zef. Questa è l'attrazione esercitata su uno specifico elettrone dal nucleo, tenendo conto di eventuali repulsioni elettrone-elettrone. Per l'idrogeno esiste un solo elettrone, quindi la carica nucleare (Z) e la carica nucleare effettiva (Zeff) sono uguali. Per tutti gli altri atomi, gli elettroni interni schermano parzialmente gli elettroni esterni dall'attrazione del nucleo, e quindi:

<img alt="Eq1" src="https://www.jove.com/files/core/CDNSource/lessons/985/985_Equations-01.png" style="height: 40px;" />

La penetrazione orbitale descrive la capacità di un elettrone di essere più vicino al nucleo. Gli elettroni nell'orbitale s possono avvicinarsi al nucleo e avere una capacità più penetrante. La densità di probabilità per un orbitale s sferico è diversa da zero nel nucleo. Diversi sottogusci hanno diversi orientamenti spaziali. A causa dell'orbitale a forma di manubrio, l'elettrone p penetra molto meno. La sua funzione d'onda ha un nodo passante per il nucleo, dove la probabilità di trovare l'elettrone è zero. Pertanto, un elettrone orbitale s è legato più strettamente al nucleo e ha un'energia inferiore rispetto all'elettrone p. Un elettrone d ha una penetrazione ancora più bassa e un'energia più elevata di un elettrone orbitale p.

Per vari gusci e sottogusci, l'andamento del potere di penetrazione di un elettrone può essere rappresentato come segue

<img alt="Eq2" src="https://www.jove.com/files/core/CDNSource/lessons/985/985_Equations-02.png" style="height: 40px;" />

L'effetto di schermatura e penetrazione è ampio e un elettrone 4s può avere un'energia inferiore rispetto a un elettrone 3d.

Questo testo è adattato da  <a href="https://openstax.org/books/chemistry-2e/pages/6-4-electronic-structure-of-atoms-electron-configurations">Openstax, Chemistry 2e, Section 6.4: Electronic Structure of Atoms (Electron Configurations).</a>

Le energie degli orbitali atomici

In un atomo, gli elettroni caricati negativamente sono attratti dal nucleo caricato positivamente. In un atomo multielettronico, le repulsioni elettrone-elettrone sono ...

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Struttura elettronica degli atomi

KEY TERMS AND CONCEPTS

TERMINI E CONCETTI CHIAVE

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...

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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

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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Shielding and Orbital Penetration

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