Name: electron filling order in atomic orbitals
Uploaded: 2020-09-03T13:38:28.000Z
Duration: 3 min 2 s
Channel: JoVE

K12 Standards

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Chemistry

Matter and its Interactions

PS.Chm.1.3

The Aufbau Principle and Hund's Rule

Bedenk dat atomaire orbitalen verschillende energieën hebben en elk twee elektronen kunnen herbergen. Het aufbau-principe dicteert de verdeling van elektronen over de subschillen van het atoom, en Hund's regel van maximale multipliciteit verklaart het vullen van orbitalen in de subschillen. Het aufbau-principe stelt dat elektronen in de grondtoestand atomaire orbitalen vullen van de laagste tot de hoogste energie om de configuratie met de laagste energie te bereiken.Hoewel de energie over het algemeen toeneemt met het aantal schillen, leidt de grotere penetratie van s-orbitalen er vaak toe dat de 4s-en 5s-orbitalen lagere energieën hebben dan de 3d-en 4d-orbitalen. De volgorde kan worden onthouden met behulp van diagrammen zoals deze, waarbij het pad van de pijl de volgorde onthult waarin elektronen aan orbitalen worden toegewezen. Overweeg om de elektronenconfiguratie voor koolstof te schrijven een element met atoomnummer zes.Zeker, de 1s-orbitaal, die de laagste energie heeft, moet vóór de 2s-orbitaal worden gevuld. Elke orbitaal kan maximaal twee elektronen bevatten. Het vijfde elektron gaat een subschil van 2p binnen.Maar welke van de drie 2p-orbitalen vult het? Nou! De orbitalen in elke subschil worden verondersteld gedegenereerd te zijn, wat betekent dat ze dezelfde energie hebben.Het vijfde elektron kan dus elk van de drie gedegenereerde 2p-orbitalen binnengaan. Hoe zit het met het zesde elektron? Komt het de 2p-orbitaal binnen met een elektron of een van de lege 2p-orbitalen?Volgens Hund's regel van maximale multipliciteit bezetten elektronen afzonderlijk alle orbitalen van een bepaald energieniveau voordat ze beginnen te paren, en de ongepaarde elektronen kunnen geen tegengestelde spins hebben:hun spins moeten parallel zijn. Daarom moeten voor koolstof de 2 2p-elektronen twee verschillende orbitalen bezetten en parallelle spins hebben. Op deze manier kunnen elektronen zich over een groter gebied verspreiden.Dit vermindert hun afscherming van elkaar, waardoor de energie van het atoom wordt geminimaliseerd. Voor stikstof is elk van de drie 2p-orbitalen afzonderlijk bezet. Voor zuurstof, zodra de gedegenereerde 2p-orbitalen afzonderlijk zijn gevuld, moet het laatste elektron paren met een ander 2p-elektron.Het atoom heeft twee ongepaarde elektronen. De elektronenconfiguratie van neon laat zien dat dit de buitenste schil is gevuld tot zijn maximale capaciteit van acht elektronen. Neon heeft twee elektronen, kernelektronen genoemd, in de binnenste schil, en acht elektronen, valentie-elektronen genoemd, in de buitenste schil.

het aufbau-principe en de regel van hund

Om de elektronenconfiguratie voor een bepaald atoom te bepalen, kunnen we de structuren in de volgorde van atoomnummers bouwen. Beginnend met waterstof, en doorgaand over de perioden van het periodiek systeem, voegen we één proton tegelijk toe aan de atoomkern en één elektron aan de juiste subschil totdat we de elektronenconfiguraties van alle elementen hebben beschreven. Deze procedure wordt het aufbau-principe genoemd, van het Duitse woord aufbau ("opbouwen"). Elk toegevoegd elektron bevindt zich in de subshell van de laagst beschikbare energie, onderworpen aan de beperkingen die worden opgelegd door de toegestane kwantumgetallen volgens het Pauli-uitsluitingsprincipe. Elektronen komen pas in subshells met hogere energie terecht nadat subshells met lagere energie volledig zijn gevuld. Figuur 1 illustreert de traditionele manier om de vulvolgorde voor atomaire orbitalen te onthouden. <img alt="Figure1" src="https://www.jove.com/files/core/CDNSource/lessons/986/986Image_1.png" style="width: 700px;"> Figuur 1 Dit diagram toont de energievolgorde voor atomaire orbitalen en is nuttig voor het afleiden van elektronenconfiguraties in de grondtoestand. Overweeg om de elektronenconfiguratie voor koolstof te schrijven, een element met atoomnummer zes. Vier elektronen vullen de 1s- en 2s-orbitalen. De overige twee elektronen bezetten de 2p-subshell. We hebben nu de keuze om één van de 2p-orbitalen te vullen en de elektronen te paren, of om de elektronen ongepaard te laten in twee verschillende, maar gedegenereerde p-orbitalen. De orbitalen worden gevuld zoals beschreven door de regel van Hund: de configuratie met de laagste energie voor een atoom met elektronen binnen een reeks gedegenereerde orbitalen is die met het maximale aantal ongepaarde elektronen. De twee elektronen in de koolstof 2p-orbitalen hebben dus identieke n-, l- en ms-kwantumgetallen en verschillen in hun ml-kwantumgetal (in overeenstemming met het Pauli-uitsluitingsprincipe). Het orbitale diagram voor koolstof, met een elektronenconfiguratie van 1s22s21p2 is: <img alt="Figure2" src="https://www.jove.com/files/core/CDNSource/lessons/986/986Image_2.png" style="width: 300px;"> Stikstof (atoomnummer 7) vult de 1s- en 2s-subschalen en heeft één elektron in elk van de drie 2p-orbitalen, in overeenstemming met de regel van Hund. Deze drie elektronen hebben ongepaarde spins. Zuurstof (atoomnummer 8) heeft een paar elektronen in een van de 2p-orbitalen (de elektronen hebben tegengestelde spins) en een enkel elektron in elk van de andere twee. Fluor (atoomnummer 9) heeft slechts één 2p-orbitaal met daarin een ongepaard elektron. Alle elektronen in het edelgas-neon (atoomnummer 10) zijn gepaard en alle orbitalen in de n = 1 en de n = 2 schillen zijn gevuld.

Het Aufbau-principe en de regel van Hund

To determine the electron configuration for any particular atom, we can build the structures in the order of atomic numbers. Beginning with hydrogen, and ...

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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 &quot;to build up.&quot; It says that each added electron enters the lowest-energy subshell available. The Pauli exclusion principle still applies, so electrons must follow the allowed quantum numbers. Higher-energy subshells do not fill until lower-energy subshells reach capacity.

A common filling-order diagram helps show the energy sequence of the atomic orbitals. Carbon is a useful example because it has atomic number 6. Four electrons fill the 1s and 2s orbitals, and the remaining two electrons go into the 2p subshell. At that point, the electrons can either pair in one 2p orbital or stay unpaired in two different p orbitals.

Hund's rule gives the lower-energy choice for degenerate orbitals, which are orbitals with the same energy. The rule states that electrons in a set of degenerate orbitals occupy the arrangement with the maximum number of unpaired electrons. In carbon, the two 2p electrons have the same n, l, and m_s quantum numbers, and they differ in m_l, which matches the Pauli exclusion principle.

Nitrogen, with atomic number 7, fills the 1s and 2s subshells and places one electron in each of the three 2p orbitals. These three electrons are unpaired. Oxygen, atomic number 8, has one pair of electrons in one 2p orbital and one single electron in each of the other two. Fluorine, atomic number 9, has only one 2p orbital with an unpaired electron.

Neon, atomic number 10, has all paired electrons. Its n = 1 and n = 2 shells are completely filled. This pattern shows how electron configurations build step by step across the first ten elements.

Electron Filling Order in Atomic Orbitals

electron-filling-order-in-atomic-orbitals

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atomic orbitals

Electron Regions in Atomic Orbitals

Electron regions in atomic orbitals show where an electron is most likely to be found in an atom. These three-dimensional regions are often drawn as boundary surfaces around the densest part of the cloud. The shape of an orbital depends on the angular momentum quantum number, l. This value can range from 0 to n - 1, where n is the principal quantum number. For n = 1, only l = 0 is allowed. For n = 2, both l = 0 and l = 1 are allowed. Orbitals with the same value of l form a subshell. The l =...

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the aufbau principle and hund's rule

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 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 electron configurations build by atomic number, why lower-energy orbitals fill first, and how Hund's rule shapes carbon, nitrogen, oxygen, fluorine, and neon.

Electron Filling Order in Atomic Orbitals | JoVE K12

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Analyze the structure of atoms and isotopes.

Electron Filling Order in Atomic Orbitals

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