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

K12 Standards

north-carolina-k-12-science-standard-course-of-study

NC K-12 Science Standard Course of Study

9-12

Chemistry

Matter and its Interactions

PS.Chm.1.3

The Aufbau Principle and Hund's Rule

원자 궤도는 서로 다른 에너지를 가지고 있고 두 개의 전자를 수용할 수 있다는 것을 상기하십시오. 쌓음 원리는 원자의 부껍질에서 전자의 분배를 나타내고, 훈드의 규칙은 부껍질에서 궤도 채우기를 설명합니다. 쌓음 원리는 바닥 상태에서 전자가 최저 에너지 구성을 달성하기 위해 에너지가 낮은 오비탈부터 차례대로 채워진다고 설명합니다.

일반적으로 에너지는 껍질 수에 따라 증가하지만, s 궤도의 침투가 크므로 4-s와 5-s 궤도가 각각 3-d와 4-d 궤도보다 낮은 에너지를 갖게 되는 경우가 많습니다. 순서는 이와 같은 도표를 사용하여 기억할 수 있습니다. 여기서 화살표의 경로는 전자가 궤도에 할당되는 순서를 나타냅니다.

원자 번호 6의 원소인 탄소에 대한 전자 구성을 작성해보십시오. 분명히 에너지가 가장 낮은 1s 궤도는 2s 궤도보다 먼저 채워져야 합니다. 매 궤도에는 최대 두 개의 전자가 들어갈 수 있습니다.

다섯 번째 전자는 2p 하위 껍질로 들어갑니다. 그러나 세 개의 2p 궤도에서 어느 것을 채웁니까? 흠~임의의 하위 껍질의 궤도는 퇴화되어 있는데, 같은 에너지를 가지고 있다는 의미입니다.

따라서 다섯 번째 전자는 세 개의 퇴하된 2p 궤도 중 어느 하나에도 들어갈 수 있습니다. 여섯 번째 전자는? 전자가 있는 2p 궤도 안으로 들어갑니까, 아니면 비어 있는 2p 궤도 중 하나로 들어갑니까?

훈트의 규칙에 따르면, 전자는 짝을 짓기 전에 주어진 에너지 수준의 모든 궤도를 단독으로 차지하며, 짝을 짓지 못한 전자는 반대 스핀을 가질 수 없습니다. 즉, 스핀은 평행이어야 합니다. 따라서 탄소의 경우 두 개의 2p 전자가 서로 다른 궤도를 차지해야 하며 평행 스핀을 가져야 합니다.

이렇게 하면 전자는 더 넓은 영역으로 퍼질 수 있습니다. 이것은 서로에 대한 차폐를 감소시켜 원자의 에너지를 최소화합니다. 질소의 경우, 3개의 2p 궤도가 각각 단독으로 점유됩니다.

산소의 경우, 일단 퇴화된 2-p 궤도를 하나씩 채우면, 마지막 전자는 다른 2-p 전자와 짝을 이루어야 합니다. 원자는 두 개의 짝을 짓지 못한 전자를 가지고 있습니다. 네온의 전자 구성은 가장 바깥쪽 껍질이 최대 8개의 용량으로 전자가 채워져 있음을 나타냅니다.

네온은 속껍질에 있는 핵심부 전자라고 불리는 두 개의 전자와 가장 바깥쪽 껍질에 있는 원자가 전자라고 불리는 여덟 개의 전자를 가지고 있습니다.

aufbau 원리와 hund의 법칙

특정 원자의 전자 구성을 결정하기 위해 원자 번호 순서대로 구조를 만들 수 있습니다. 수소부터 시작하여 주기율표의 주기에 걸쳐 계속해서 모든 원소의 전자 구성을 설명할 때까지 핵에 한 번에 하나의 양성자를 추가하고 적절한 부 껍질에 하나의 전자를 추가합니다. 이 절차를 독일어 단어 aufbau(“쌓다”)에서 유래한 aufbau 원리라고 합니다. 추가된 각 전자는 파울리 배제 원리에 따라 허용된 양자수에 의해 부과된 제한에 따라 사용 가능한 가장 낮은 에너지의 부 껍질을 차지합니다. 전자는 낮은 에너지 부 껍질이 용량까지 채워진 후에만 더 높은 에너지 부 껍질로 들어갑니다. 그림 1은 원자 궤도의 채우기 순서를 기억하는 전통적인 방법을 보여줍니다.

<img alt="Figure1" src="https://www.jove.com/files/core/CDNSource/lessons/986/986Image_1.png" style="width: 700px;" />

Figure 1 이 다이어그램은 원자 궤도의 에너지 순서를 나타내며 바닥 상태 전자 구성을 도출하는 데 유용합니다.

원자번호 6번 원소인 탄소의 전자 배치를 작성해 보세요. 4개의 전자가 1s와 2s 오비탈을 채웁니다. 나머지 2개의 전자는 2p 부껍질을 차지합니다. 이제 우리는 2p 오비탈 중 하나를 채우고 전자를 쌍으로 만들거나 두 개의 서로 다르지만 축퇴된 p 오비탈에 짝을 이루지 않은 전자를 남겨 둘 수 있습니다. 궤도는 Hund의 규칙에 설명된 대로 채워집니다. 축퇴 궤도 집합 내에 전자가 있는 원자의 가장 낮은 에너지 구성은 짝을 이루지 않은 전자의 최대 수를 갖는 것입니다. 따라서 탄소 2p 오비탈에 있는 두 전자는 동일한 n, l, ms 양자수를 가지며 ml 양자수는 다릅니다(Pauli 배제 원리에 따라). 1s^22s^21p^2의 전자 구성을 갖는 탄소의 궤도 다이어그램은 다음과 같습니다.

<img alt="Figure2" src="https://www.jove.com/files/core/CDNSource/lessons/986/986Image_2.png" style="width: 300px;" />

질소(원자번호 7번)는 훈트의 법칙에 따라 1s와 2s 부껍질을 채우고 3개의 2p 오비탈 각각에 하나의 전자를 가지고 있습니다. 이 세 개의 전자는 짝을 이루지 않은 스핀을 가지고 있습니다. 산소(원자 번호 8)는 2p 오비탈 중 하나에 전자 쌍(전자는 반대 스핀을 가짐)을 갖고 나머지 두 오비탈에는 각각 전자 1개를 가지고 있습니다. 플루오린.(원자번호 9번)는 짝을 이루지 않은 전자를 포함하는 2p 오비탈을 하나만 가지고 있습니다. 희가스 네온(원자번호 10)의 전자는 모두 쌍을 이루며 n=1 껍질과 n=2 껍질의 오비탈이 모두 채워집니다.

이 문서는 에서 발췌되었습니다 <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.</a>

Aufbau 원리와 Hund의 법칙

특정 원자에 대한 전자 구성을 결정하기 위해 원자 번호 순서로 구조를 구축할 수 있습니다. 수소로 시작하여 ...

Education

교육

JoVE Core

화학

하나

Electronic Structure of Atoms

원자의 전자 구조

KEY TERMS AND CONCEPTS

주요 용어 및 개념

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

the wave nature of light

Measuring Light with Wavelength and Frequency

Light can be described as an electromagnetic wave, and its behavior helps explain many patterns seen in science. For centuries, scientists have tried to understand light. Newton used lenses and prisms to show that white light is made of the colors of the rainbow mixed together. He explained light as tiny particles moving at high speed. Other scientists, including Christiaan Huygens, showed that reflection and refraction could also be explained by treating light as a wave. At the time, people...

the electromagnetic spectrum

Electromagnetic Waves: Uses Across the Spectrum

The electromagnetic spectrum includes all electromagnetic radiation arranged by frequency and wavelength. Visible light is only a small part of it, and each color of visible light has its own frequency and wavelength. Different parts of the spectrum also use different units. Radio waves are usually measured in MHz, while the visible range is often measured in nm or angstroms. The spectrum is organized from radio waves to gamma rays as frequency increases and wavelength decreases. Radio waves...

interference and diffraction

Wave Patterns from Light and Slits

Light can form wave patterns through interference and diffraction. When two electromagnetic waves meet with their peaks and troughs lined up, they combine into a larger wave. This is constructive interference, and the waves are in phase. When a peak of one wave meets a trough of the other, the waves reduce each other. This is destructive interference, and the result has a much lower amplitude. Diffraction is another wave behavior. When light passes through a slit that is about the same size as...

photoelectric effect

Why Light Ejects Electrons from Metals

Light can eject electrons from a metal surface when it has the right frequency. This process is called the photoelectric effect. Each metal has its own threshold frequency, which is the minimum light frequency needed to free electrons. Light below the threshold frequency cannot cause emission, even if it is very bright. When the light frequency is above the threshold, electrons are ejected. In that case, the number of electrons released is directly proportional to the intensity, or brightness,...

the bohr model

Bohr’s Quantized Atomic Orbits

Bohr’s quantized atomic orbits explain why hydrogen atoms can stay stable. Early in the twentieth century, Ernest Rutherford and his colleagues described atoms as tiny dense nuclei with lighter electrons moving around them. This picture was called the planetary model because it compared the atom to a miniature solar system, with electrons orbiting the nucleus like planets around the sun. The simplest atom is hydrogen. It has one proton in the nucleus and one electron moving around it. In the...

emission spectra

Line Spectra from Heated Gases

Line spectra are produced when heated gases give off light at specific wavelengths. In contrast, hot solids, liquids, or condensed gases radiate excess energy as a continuous spectrum. That spectrum contains an unbroken range of wavelengths. A gas at low partial pressure can produce a line spectrum when it is excited by an electric current or heated. Fluorescent light bulbs and neon signs work in this way. Each element has its own characteristic set of lines, and molecules can also produce...

the de broglie wavelength

Electron Waves and Standing Orbits

Electron waves and standing orbits explain how tiny particles can behave like waves. Large objects in the macroscopic world, like a billiard ball on a table, follow classical physics. A ball keeps moving in a straight line unless it hits another ball or a force such as friction acts on it. It also has a well-defined position, velocity, and momentum, where momentum is mass times velocity. Waves behave differently from macroscopic particles. When waves interact, they can form interference...

the uncertainty principle

Electron Position and Momentum Limits

Heisenberg’s uncertainty principle sets a limit on how precisely we can know an electron’s position and momentum at the same time. Momentum is the motion of a particle, and the two measurements affect each other. If the momentum is known more accurately, the position becomes less certain, and the reverse is also true. Heisenberg described this limit with an equation that uses Planck’s constant. The equation shows the smallest possible uncertainty for position and momentum together. It explains...

the quantum-mechanical model of an atom

Electron Orbitals in Quantum Mechanics

Quantum mechanics explains how electrons behave in atoms using waves and probability. Erwin Schrödinger built on de Broglie’s idea that an electron in hydrogen could act like a circular standing wave instead of moving in fixed circular paths. He then derived the Schrödinger equation, which became a major step in modern atomic theory. When Schrödinger applied his equation to hydrogen-like atoms, it matched Bohr’s energy expression and the Rydberg formula for hydrogen spectra. He described...

quantum numbers

Electron Orbitals and Spin

Electron orbitals are described by quantum numbers that show where an electron is most likely to be found around the nucleus. These numbers come from the quantum mechanical model, which is based on solutions of the Schrödinger equation. They also help explain why electron energy comes in set values instead of changing smoothly. The principal quantum number, n, labels the energy level. It can be 1, 2, 3, and so on, and higher n values usually mean higher electron energy. This number also tells...

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 =...

the pauli exclusion principle

Electron Spin and Orbital Pairing

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

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 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

electron-configuration-of-multielectron-atoms

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 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

chemistry

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Louisiana Student Standards for Science

High School

Matter And Its Interactions

HS-PS1-1

HS-PS1-2

ohios-learning-standards-and-model-curriculum

Ohio’s Learning Standards and Model Curriculum

Structure And Properties Of Matter

Atomic structure

Tennessee Academic Standards

Matter and Its Interactions

CHEM1.PS1.12

Virginia  Standards of Learning

HIgh School

CH.2

CH.2.e

Physical Science

PS.2

PS.2.a

Arizona Science Standards

HS.P1

HS.P1U1.1

Texas Essential Knowledge and Skills

CHEM.6

CHEM.6.B

Alabama Course of Study Science

9-12.C.1.1.c

Next Generation Science Standards

Florida Department of Education

SC.912.P.8

SC.912.P.8.5

SC.912.P.10

SC.912.P.10.9

new-york-state-science-learning-standards

New York State Science Learning Standards

HS. Structure and Properties of Matter

HS-PS1-1.

HS. Chemical Reactions

HS-PS1-7.

Georgia Standards of Excellence

SC1.

electron-configuration-and-chemical-properties

Electron Configuration and Chemical Properties

Scientific Laws and Theories

Scientific Laws

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The Scientific Method

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Classifying Matter by State

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Classifying Matter by Composition

Matter: Pure Substances and Mixtures
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The Atomic Theory of Matter

The earliest recorded discussion of the basic structure of matter comes from ancient Greek philosophers. Leucippus and Democritus argued that all matter ...

Subatomic Particles

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Elements: Chemical Symbols and Isotopes

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Use models to explain how the scientific understanding of atomic structure has evolved.

PS.Chm.1.1

Radioactivity and Nuclear Equations

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Types of Radioactivity

The most common types of radioactivity are &#945; decay, &#946; decay, &#947; decay, neutron emission, and electron capture.
Alpha (&#945;) decay is the ...

Nuclear Stability

Protons and neutrons, collectively called nucleons, are packed together tightly in a nucleus. With a radius of about 10&#8722;15 meters, a nucleus is ...

Nuclear Binding Energy

The difference between the calculated and experimentally measured masses is known as the mass defect of the atom. In the case of helium-4, the mass defect ...

Radioactive Decay and Radiometric Dating

Radioactivity is a spontaneous disintegration of an unstable nuclide and is a random process, as all the nuclei in the sample do not decay simultaneously. ...

Nuclear Fission

Many heavier elements with smaller binding energies per nucleon can decompose into more stable elements that have intermediate mass numbers and larger ...

Nuclear Power

Controlled nuclear fission reactions are used to generate electricity. Any nuclear reactor that produces power via the fission of uranium or plutonium by ...

Nuclear Fusion

The process of converting very light nuclei into heavier nuclei is also accompanied by the conversion of mass into large amounts of energy, a process ...

Nuclear Transmutation

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Use models to compare nuclear reactions including alpha decay, beta decay and gamma decay; nuclear fusion and nuclear fission.

PS.Chm.1.2

Atomic Orbitals

An atomic orbital represents the three-dimensional regions in an atom where an electron has the highest probability to reside. The radial distribution ...

The Pauli Exclusion Principle

The arrangement of electrons in the orbitals of an atom is called its electron configuration. We describe an electron configuration with a symbol that ...

The Energies of Atomic Orbitals

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

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

Electron Configuration of Multielectron Atoms

The alkali metal sodium (atomic number 11) has one more electron than the neon atom. This electron must go into the lowest-energy subshell available, the ...

Use models to explain how electrons are distributed in atoms.

Periodic Classification of the Elements

The periodic table arranges atoms based on increasing atomic number so that elements with the same chemical properties recur periodically. When their ...

The Periodic Table

As early chemists discovered more elements, they realized that various elements could be grouped by their similar chemical behaviors. One such grouping ...

Atomic Mass

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Molar Mass

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Use the Periodic Table as a model to predict the relative properties of elements based on the pattern of valence electrons in the outermost energy levels of atoms.

PS.Chm.2.1

atomic-radii-and-effective-nuclear-charge

Atomic Radii and Effective Nuclear Charge

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Ionic Radii

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Ionization Energy

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Electron Affinity

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Construct an explanation to infer the atomic size, reactivity, electronegativity, and ionization energy of an element based on its position in the Periodic Table.

PS.Chm.2.2

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Lewis Symbols and the Octet Rule

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Ionic Bonding and Electron Transfer

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The Born-Haber Cycle

Lattice Energy&#xA0;

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Covalent Bonding and Lewis Structures

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Electronegativity

Whether a bond is nonpolar or polar covalent is determined by a property of the bonding atoms called electronegativity.&#xA0;

Electronegativity values of ...

Analyze and interpret data to explain the mechanisms and properties of the two main types of intramolecular (ionic and covalent) bonds.

PS.Chm.3.1

Analyze the structure of atoms and isotopes.

Electron Filling Order in Atomic Orbitals

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