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

K12 Standards

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High School

Chemistry and Physics

IPC.7

IPC.7.E

The Energies of Atomic Orbitals

Atoomorbitalen hebben verschillende energieën, zoals gerationaliseerd door Coulomb-interacties, het afschermende effect en orbitale penetratie. De wet van Coulomb geeft aan dat de aantrekkingskracht of afstotende kracht tussen twee geladen deeltjes een omgekeerd-kwadratische relatie heeft met de afstand ertussen. Atomaire orbitale afmetingen nemen toe met het aantal schillen en elektronen worden afgestoten uit de ruimte die wordt ingenomen door orbitalen met een lagere schil.De wet van Coulomb suggereert dus dat naarmate het aantal schillen toeneemt, de elektronen minder aangetrokken worden tot de kern, wat overeenkomt met hogere orbitale energieën. Bovendien hebben elektronen die zich op ongeveer dezelfde afstand van of dichter bij de kern bevinden een afschermend effect dat de aantrekking naar de kern verder vermindert. Hoe groter de afscherming, hoe minder aantrekkingskracht op de kern wordt gevoeld.Dit is een reden voor de verschillen in orbitale energieën binnen elektronenschillen. Bijvoorbeeld, 3s en 3p elektronen beschermen 3d-elektronen aanzienlijk. De effectieve kernlading die door een elektron wordt gevoeld, wordt berekend door de afschermconstante S, die afhangt van het aantal afschermende elektronen en de subschillen die ze innemen, af te trekken van het atoomnummer.De twee 1s-elektronen in lithium bijvoorbeeld, dat een atoomnummer van drie heeft, schermen het 2s-elektron af. De afschermconstante voor dat elektron wordt op basis van semi-empirische regels bepaald op 1, 7. Daarom is de effectieve nucleaire lading die door het 2s-elektron wordt gevoeld 1, 3.De vormen van orbitalen dicteren ook hun energie. Als de elektronen in een buitenste orbitaal ver kunnen bewegen in gebieden die worden ingenomen door binnenste elektronen om dicht bij de kern te zijn, zullen ze daar veel minder afgeschermd zijn. Dus de energie van die buitenste orbitaal is lager.Dit kan worden gevisualiseerd met een radiale verdelingsfunctie die de kans beschrijft om een elektron op een bepaalde afstand van de kern te vinden. Radiale verdelingsfunctiegrafieken voor de 1s, 2s en 2p subschillen laten zien dat 2s-elektronen een bescheiden kans hebben om in de buurt van de kern te zijn, terwijl 2p-elektronen meestal buiten of aan de buitenrand van het 1s-gebied blijven. De 2s-orbitaal zou daarom een groter doordringend vermogen hebben.In de derde schil dringen de 3s-elektronen het meest door en de 3d-elektronen het minst. Over het algemeen neemt de atoomomloop-energie toe met het aantal schillen en, op het niveau van de schil, van s naar f. Het penetratie-effect wordt echter zo belangrijk in de vierde en vijfde schil dat de 4s en 5s orbitalen vaak lagere relatieve energieën hebben dan respectievelijk de 3d-en 4d-orbitalen.

de energieën van atomaire orbitalen

In een atoom worden de negatief geladen elektronen aangetrokken door de positief geladen atoomkern. In een multi-elektronenatoom worden ook elektron-elektron-afstotingen waargenomen. De aantrekkende en afstotende krachten zijn afhankelijk van de afstand tussen de deeltjes, evenals van het teken en de grootte van de ladingen op de individuele deeltjes. Wanneer de ladingen op de deeltjes tegengesteld zijn, trekken ze elkaar aan. Als beide deeltjes dezelfde lading hebben, stoten ze elkaar af. Naarmate de omvang van de ladingen toeneemt, neemt de omvang van de kracht toe. Wanneer de scheiding van ladingen echter groter is, nemen de krachten af. De aantrekkingskracht tussen een elektron en zijn atoomkern is dus recht evenredig met de afstand daartussen. Als het elektron dichter bij de atoomkern is, is het nauwer aan de atoomkern gebonden; daarom hebben de elektronen in de verschillende schillen (op verschillende afstanden) verschillende energieën. Bij atomen met meerdere energieniveaus beschermen de binnenste elektronen de buitenste elektronen gedeeltelijk tegen de aantrekkingskracht van de atoomkern, als gevolg van elektron-elektronafstoting. Kernelektronen beschermen elektronen in de buitenste schillen, terwijl elektronen in dezelfde valentieschil de nucleaire aantrekkingskracht die elkaar ervaart niet zo efficiënt blokkeren. Dit kan worden verklaard met het concept van effectieve nucleaire lading, Zeff. Dit is de aantrekkingskracht die de atoomkern uitoefent op een specifiek elektron, waarbij rekening wordt gehouden met eventuele elektron-elektron-afstoting. Voor waterstof is er maar één elektron, en dus zijn de nucleaire lading (Z) en de effectieve nucleaire lading (Zeff) gelijk. Voor alle andere atomen beschermen de binnenste elektronen de buitenste elektronen gedeeltelijk tegen de aantrekkingskracht van de atoomkern, en dus: <img alt="Eq1" src="https://www.jove.com/files/core/CDNSource/lessons/985/985_Equations-01.png" style="height: 40px;"> Orbitale penetratie beschrijft het vermogen van een elektron om dichter bij de atoomkern te zijn. De elektronen in de s-orbitaal kunnen dichter bij de atoomkern komen en hebben een groter doordringend vermogen. De waarschijnlijkheidsdichtheid voor een bolvormige s-orbitaal is bij de atoomkern niet nul. Verschillende subshells hebben verschillende ruimtelijke oriëntaties. Door de haltervormige orbitaal dringt het p-elektron veel minder binnen. De golffunctie heeft een knooppunt dat door de atoomkern gaat, waar de kans om het elektron te vinden nul is. Een s-orbitaal elektron is dus nauwer aan de atoomkern gebonden en heeft een lagere energie dan het p-elektron. Een d-elektron heeft een nog lagere penetratie en hogere energie dan een p-orbitaal elektron. Voor verschillende schillen en subschillen kan de trend van het doordringend vermogen van een elektron als volgt worden weergegeven <img alt="Eq2" src="https://www.jove.com/files/core/CDNSource/lessons/985/985_Equations-02.png" style="height: 40px;"> Het effect van afscherming en penetratie is groot en een 4s-elektron kan een lagere energie hebben dan een 3d-elektron.

De energieën van atomaire orbitalen

In an atom, the negatively charged electrons are attracted to the positively charged nucleus. In a multielectron atom, electron-electron repulsions are ...

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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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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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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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