Name: shielding and orbital penetration
Uploaded: 2020-09-03T13:38:28.000Z
Duration: 3 min 21 s
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K12 Standards

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High School

Chemistry and Physics

IPC.7

IPC.7.E

The Energies of Atomic Orbitals

Les orbitales atomiques ont des énergies différentes, telles que rationalisées par les interactions de Coulomb, l'effet de blindage et la pénétration orbitale. La loi de Coulomb indique que la force d'attraction ou de répulsion entre deux particules chargées a une relation inverse-carrée avec la distance entre elles. La taille des orbitales atomiques augmente avec le nombre de shell, et les électrons sont repoussés de l'espace occupé par les orbitales inférieures.Ainsi, la loi de Coulomb suggère que lorsque le nombre de shell augmente, les électrons éprouvent moins d'attraction vers le noyau, ce qui correspond à des énergies orbitales plus élevées. En outre, les électrons qui sont à peu près à la même distance ou plus près du noyau ont un effet de blindage qui réduit encore l'attraction vers le noyau. Plus le blindage est grand, moins l'attraction vers le noyau est ressentie.C'est l'une des raisons des différences d'énergies orbitales dans les shells d'électrons. Par exemple, les électrons 3s et 3p protègent significativement les électrons 3d. La charge nucléaire effective ressentie par un électron est calculée en soustrayant du numéro atomique la constante de blindage S, qu i dépend du nombre d'électrons de blindage et des sous-shells qu'ils occupent.Par exemple, les deux électrons 1s dans le lithium, qui a un numéro atomique de trois, écran son électron 2s. La constante de blindage pour cet électron est déterminée à partir de règles semi-empiriques à 1.7. Par conséquent, la charge nucléaire effective ressentie par l'électron 2s est de 1, 3.Les formes des orbitales dictent également leur énergie. Si les électrons d'une orbitale extérieure peuvent se déplacer loin dans des zones occupées par des électrons internes pour être proches du noyau, ils y seront beaucoup moins protégés. Ainsi, l'énergie de cette orbitale externe est plus faible.Ceci peut être visualisé avec une fonction de distribution radiale décrivant la probabilité de trouver un électron à une distance donnée du noyau. Les diagrammes de fonction de distribution radiale pour les sous-shells 1s, 2s et 2p révèlent que les électrons 2s ont une probabilité modeste d'être près du noyau, alors que les électrons 2p restent principalement à l'extérieur ou au bord externe de la région 1s. On dit donc que l'orbite 2s a une plus grande capacité de pénétration.Dans la troisième shell, les électrons 3s pénètrent le plus et les électrons 3d pénètrent le moins. En général, l'énergie orbitale atomique augmente avec le nombre de shell et, au niveau du sous-coquillage, de s à f. Cependant, l'effet de pénétration devient si important dans les quatrième et cinquième shells que les orbitales 4s et 5s ont souvent des énergies relatives plus faibles que les orbitales 3d et 4d, respectivement.

les énergies des orbitales atomiques

Dans un atome, les électrons chargés négativement sont attirés par le noyau chargé positivement. Dans un atome à plusieurs électrons, on observe également des répulsions électron-électron. Les forces attractives et répulsives dépendent de la distance entre les particules, ainsi que du signe et de l'intensité des charges sur les particules individuelles. Lorsque les charges sur les particules sont opposées, elles s'attirent les unes les autres. Si les deux particules ont la même charge, elles se repoussent l'une l'autre. 
 
 À mesure que l'intensité des charges augmente, l'intensité de la force augmente. Toutefois, lorsque la séparation des charges est plus importante, les forces diminuent. Ainsi, la force d'attraction entre un électron et son noyau est directement proportionnelle à la distance qu'il y a entre eux. Si l'électron est plus proche du noyau, il se lie plus fermement au noyau ; par conséquent, les électrons sur les différentes couches (à différentes distances) ont des énergies différentes. 
 
 Pour les atomes à plusieurs niveaux d'énergie, les électrons internes protègent partiellement les électrons externes de l'attraction du noyau, en raison des répulsions électron-électron. Les électrons de cœur protègent les électrons sur les couches externes, tandis que les électrons sur la même couche de valence ne bloquent pas l'attraction nucléaire qu'ils subissent les uns les autres aussi efficacement. Cela peut être expliqué grâce au concept de la charge nucléaire effective, Zeff. Il s'agit de la force exercée par le noyau sur un électron précis, en tenant compte des répulsions électron-électron. Pour l'hydrogène, il n'y a qu'un seul électron, et la charge nucléaire (Z) et la charge nucléaire effective (Zeff) sont donc égales. Pour tous les autres atomes, les électrons internes protègent partiellement les électrons externes de l'attraction du noyau, et donc : 
 
 <img alt="Eq1" src="https://www.jovecore.com/CDNSource/lessons/985/985_Equations-01.png" style="height: 40px;" />
 
 La pénétration d'une orbitale décrit la capacité d'un électron à être plus proche du noyau. Les électrons de l'orbitale s peuvent se rapprocher du noyau et ils ont une capacité plus pénétrante. La densité de probabilité d'une orbitale s sphérique est non nulle au niveau du noyau.  Les différentes sous-couches ont des orientations spatiales différentes. En raison de l'orbitale en forme d'haltère, l'électron p pénètre beaucoup moins. Sa fonction d'onde a un nœud passant à travers le noyau, où la probabilité de trouver l'électron est zéro. Ainsi, un électron de l'orbitale s est lié plus fermement au noyau et a une énergie plus faible que l'électron p. Un électron d a une pénétration encore plus faible et une énergie plus élevée qu'un électron de l'orbitale p. Pour les diverses couches et sous-couches, la tendance de la puissance pénétrante d'un électron peut être décrite comme suit 
 
 <img alt="Eq2" src="https://www.jovecore.com/CDNSource/lessons/985/985_Equations-02.png" style="height: 40px;" />
 
 L'effet d'écran et de la pénétration est important, et un électron 4s peut avoir une énergie inférieure à celle d'un électron 3d. 
 
 Ce texte est adapté de <a href="https://openstax.org/books/chemistry-2e/pages/6-4-electronic-structure-of-atoms-electron-configurations">Openstax, Chimie 2e, Section 6.4 : Structure électronique des atomes (configurations électroniques).</a>

Les énergies des orbitales atomiques

Dans un atome, les électrons chargés négativement sont attirés par le noyau chargé positivement. Dans un atome de multiélectron, les répulsions électron-électron sont ...

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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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Electron spin and orbital pairing explain why only two electrons can fit in one atomic orbital. Electron configuration describes how electrons are arranged in an atom’s orbitals. It is written with three parts: the principal quantum number, n, the orbital type, or subshell, and a superscript that shows how many electrons are in that subshell. For example, 2p4 means four electrons in a p subshell with n = 2, and 3d8 means eight electrons in a d subshell with n = 3. The three quantum numbers...

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.

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Shielding and Orbital Penetration

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