Name: how electrons fill atomic orbitals
Uploaded: 2023-04-30T18:28:31.000Z
Duration: 2 min 46 s
Channel: JoVE

K12 Standards

Louisiana Student Standards for Science

High School

Chemistry

Matter And Its Interactions

HS-PS1-1

Electron Configurations

يمثل تكوين الإلكترون للذرة توزيع الإلكترونات بين مداراتها الذرية. يمكن توسيع مبدأ استبعاد باولي ، وقاعدة هوند للحد الأقصى للتعددية ، ومبدأ أوفباو لتصور تكوين الإلكترون لأي عنصر.
ينص مبدأ Aufbau على أنه في الحالة الأرضية ، تملأ المدارات الذرية بترتيب متزايد من الطاقة. يتم ترشيد الطاقات النسبية للمدارات الذرية من خلال تفاعلات كولوم ، وتأثير التدريع ، والاختراق المداري.
ضع في اعتبارك تكوين الإلكترون للكربون - وهو عنصر ذو رقم ذري ستة وبالتالي محايد بستة إلكترونات. يتم ملء المدار 1s ، الذي يحتوي على أقل طاقة ، أولا.
وفقا لمبدأ استبعاد باولي ، لا يمكن أن يكون لإلكترونين في الذرة نفس المجموعة المكونة من أربعة أرقام كمومية. نظرا لأن الإلكترونات في نفس المدار لها نفس الأرقام الكمومية الرئيسية والسمت والمغناطيسية ، يجب أن يكون لها أرقام كمومية دورانية مختلفة.
نظرا لأن الرقم الكمي للدوران له قيمتان محتملتان فقط ، يمكن للمدار أن يستوعب إلكترونين فقط ، مع دورات متعاكسة.
على الرغم من أن الطاقة تزداد مع عدد الغلاف ، إلا أن الاختراق الأكبر للمدارات s يقلل من طاقة المدارات s بالنسبة لتلك الموجودة في المدارات p.
لذلك ، يشغل الإلكترونات التاليان مدار 2s ويدخل الخامس الغلاف الفرعي 2p. نظرا لأنه يفترض أن المدارات داخل غلاف فرعي متدهورة ، يمكن للإلكترون الخامس أن يدخل أيا من المدارات الثلاثة المتدهورة 2ص.
يتبع الإلكترون السادس قاعدة هوند المتمثلة في التعددية القصوى ويحتل مفردا مدارا متدهورا بدلا من الاقتران. ومن ثم ، بالنسبة للكربون ، يشغل الإلكترونين 2p مدارين مختلفين ولهما دورات متوازية.
يكشف مخطط تكوين الإلكترون أن الكربون يحتوي على إلكترونين - يسمى الإلكترونات الأساسية - في الغلاف الداخلي وأربعة إلكترونات - تسمى إلكترونات التكافؤ - في الأغلفة الخارجية.
يمكن تذكر ترتيب المدارات الذرية بالنسبة لطاقاتها باستخدام مثل هذه المخططات ، حيث يكشف مسار السهم عن التسلسل الذي يتم فيه تعيين الإلكترونات للمدارات.
ومع ذلك ، يصبح هذا التقدم أكثر تعقيدا مع إدخال المدارات d و f ويصبح التسلسل أقل قابلية للتنبؤ بالعناصر الانتقالية واللانثانيدات والأكتينيدات.

توزيعات الإلكترونات

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يمكن تحديد توزيعات الإلكترونات والمخططات المدارية من خلال تطبيق مبدأ أوفباو (يشغل كل إلكترون مضاف أقل مستوى طاقة متاح من المدارات الفرعية)، ومبدأ باولي للاستبعاد (لا يمكن أن يكون لإلكترونين نفس مجموعة الأعداد الكمية الأربعة)، وقاعدة هوند للتعددية القصوى (كلما كان ذلك ممكنًا، تحتفظ الإلكترونات بالدوران غير المزدوج في المدارات المتكافئة).

تحدد الطاقات النسبية للمدارات الفرعية الترتيب الذي يتم به ملء المدارات الذرية ((1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p,، وهكذا). بالنسبة للأغلفة والأغلفة الفرعية المختلفة، يمكن وصف اتجاه قوة اختراق الإلكترون على النحو التالي:

1s > 2s > 2p > 3s > 3p > 4s > 3d> 4p > 5s > 4d > 5p > 6s > f.…

تأثير التدريع والاختراق المداري كبير، وقد يكون للإلكترون 4s طاقة أقل من الإلكترون 3d.

الإلكترونات الموجودة في المدارات الخارجية، والتي تسمى إلكترونات التكافؤ، هي المسؤولة عن معظم السلوك الكيميائي للعناصر. في الجدول الدوري، عادة ما توجد العناصر ذات التوزيعات الإلكترونية المتماثلة داخل نفس المجموعة.

هناك بعض الاستثناءات لقاعدة الملء المتوقعة، خاصة عندما يمكن تكوين مدارات نصف ممتلئة أو ممتلئة بالكامل. في حالة Cr وCu، يبدو أن الأغلفة الفرعية نصف المملوءة والمملوءة بالكامل تمثل ظروف الاستقرار المفضل. هذا الاستقرار هو أن الإلكترون ينتقل من 4s إلى المدار 3d للحصول على استقرار إضافي لمدار فرعي 3d نصف ممتلئ (في Cr) أو مدار فرعي 3d ممتلئ (في Cu). تحدث استثناءات أخرى أيضًا. على سبيل المثال، من المتوقع أن يكون للنيوبيوم (Nb، العدد الذري 41) التكوين الإلكتروني[Kr]5s24d3 ومع ذلك، من الناحية التجريبية، فإن التكوين الإلكتروني لحالة الاستقرار هو في الواقع [Kr]5s14d4. يمكننا تبرير هذه الملاحظة بالقول إن قوى التنافر بين الإلكترونات الناتج عن اقتران الإلكترونات في المدار 5s أكبر من فجوة الطاقة بين المدارات 5s و4d.

<a href="https://openstax.org/books/chemistry-2e/pages/6-4-electronic-structure-of-atoms-electron-configurations">This text is adapted from  Openstax, Chemistry 2e, Section 6.4: Electronic Structure of Atoms</a>

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يمكن تحديد تكوينات الإلكترون والمخططات المدارية من خلال تطبيق مبدأ Aufbau (كل إلكترون مضاف يحتل الغلاف الفرعي لأدنى ...

Education

JoVE Core

Organic Chemistry

الكيمياء العضوية

Unit 1

الوحدة الأولى

Covalent Bonding and Structure

الترابط التساهمي والبنية

KEY TERMS AND CONCEPTS

المصطلحات والمفاهيم الأساسية

how electrons fill atomic orbitals

Electron filling in atoms follows a set order based on orbital energy. The Aufbau principle says each new electron enters the lowest-energy subshell available. The Pauli exclusion principle says no two electrons can share the same set of four quantum numbers. Hund’s rule of maximum multiplicity says electrons stay unpaired in equal-energy orbitals when possible.

The order of filling follows the relative energies of the subshells. A common sequence is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, and so on. This order is shaped by shielding and orbital penetration, which affect how strongly an electron feels the nucleus. A 4s electron may have lower energy than a 3d electron because of these effects.

Valence electrons are the electrons in the outermost orbitals. They control most of the chemical behavior of an element. In the periodic table, elements with similar valence electron configurations usually appear in the same group. That pattern helps explain why group members often behave in similar ways.

The predicted filling order does not always match the real ground-state configuration. Some exceptions appear when half-filled or fully filled subshells are more stable. Chromium and copper are classic examples. In chromium, an electron shifts from 4s into 3d to make a half-filled 3d subshell. In copper, the shift creates a filled 3d subshell.

Other exceptions also occur. Niobium, with atomic number 41, is predicted to have the configuration [Kr]5s2 4d3. Experimentally, its ground-state electron configuration is [Kr]5s1 4d4. This can be explained by electron-electron repulsion in the 5s orbital. In this case, that repulsion is larger than the energy gap between the 5s and 4d orbitals.

How Electrons Fill Atomic Orbitals

what is organic chemistry?

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Atomic Orbitals and Quantum Numbers

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electron configurations

Electron filling in atoms follows a set order based on orbital energy. The Aufbau principle says each new electron enters the lowest-energy subshell available. The Pauli exclusion principle says no two electrons can share the same set of four quantum numbers. Hund’s rule of maximum multiplicity says electrons stay unpaired in equal-energy orbitals when possible. The order of filling follows the relative energies of the subshells. A common sequence is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, and so on.

chemical bonds

Chemical Bonds: Types and Strength

Chemical bonds help atoms reach a full valence shell like the nearest noble gas. The main bond types are ionic, covalent, and metallic. Bond energy and bond length also help explain how strong a bond is. An ionic bond forms from electrostatic attraction between cations and anions. In many cases, one atom transfers electrons to another atom, and the ions form from that transfer. Ionic bonds have no defined direction because the attraction spreads evenly through three-dimensional space. A...

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Electronegativity in Polar Covalent Bonds

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Resonance in Nitrite and Carbonate Ions

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Covalent Bond Overlap and Orbital Hybridization

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mo theory and covalent bonding

Bond Order and Molecular Orbitals

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intermolecular forces and physical properties

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introduction to functional groups

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Learn how Aufbau, Pauli exclusion, and Hund’s rule guide orbital filling, why valence electrons matter, and why chromium, copper, and niobium are exceptions.

How Electrons Fill Atomic Orbitals | JoVE K12

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HS-PS1-2

Tennessee Academic Standards

Matter and Its Interactions

CHEM1.PS1.12

Virginia  Standards of Learning

HIgh School

CH.2

CH.2.d

Minnesota Academic Standards in Science

9-12

3.1.1 9C.2.1.1.1

Next Generation Science Standards

Physical Science

Florida Department of Education

SC.912.P.8

SC.912.P.8.5

SC.912.P.8.12

south-carolina-college-and-career-ready-science-standards

South Carolina College and Career Ready Science Standards

C-PS1-1

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Electric Fields

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Intermolecular Forces

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Van der Waals Interactions

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Chemical Bonds


Atoms participate in a chemical bond formation&#160;to acquire a completed valence-shell electron configuration similar to that of the noble gas nearest ...

Electron Behavior

Electrons are negatively charged subatomic particles attracted to and orbit around the positively-charged nucleus of an atom. They reside in spaces ...

Subatomic Particles

Dalton was only partially correct about the particles that make up matter. All matter is composed of atoms, and atoms are composed of three smaller ...

Ions and Ionic Charges

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Use the periodic table as a model to predict the relative properties of elements based on the patterns of electrons in the outermost energy level and the composition of the nucleus of atoms.  Clarification Statement: Chemistry: Examples of properties that could be predicted from patterns could include reactivity of metals, types of bonds formed, numbers of bonds formed, atomic radius, atomic mass, or reactions with oxygen. Emphasis is on main group elements and qualitative understanding of the relative trends of ionization energy and electronegativity.

Properties of Transition Metals

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The Atomic Theory of Matter

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Chemical Equations

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The Collision Theory
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Stoichiometry, Product Yield, and Limiting Reactants

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Physical Properties of Matter

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Dynamic Equilibrium

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Construct and revise an explanation for the outcome of a simple chemical reaction based on the outermost electron states of atoms, trends in the periodic table, and knowledge of the patterns of chemical properties.  Clarification Statement: Chemistry: Examples of chemical reactions could include the reaction of sodium and chlorine, carbon and oxygen, or carbon and hydrogen. Reaction classification aids in the prediction of products (e.g. synthesis, decomposition, single displacement, double displacement, and acid-base).

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Intermolecular Forces and Physical Properties



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Electric Field of Two Equal and Opposite Charges

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Electric Dipoles and Dipole Moment

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Plan and conduct an investigation to gather evidence to compare the structure of substances at the macroscale to infer the strength of electrical forces between particles.  Clarification Statement: Emphasis is on understanding the strengths of forces between particles, not on naming specific intermolecular forces (such as dipole-dipole). Examples of particles could include ions, atoms, molecules, and network solids (such as graphite). Examples of macro-properties of substances could include the melting point and boiling point, vapor pressure, and surface tension.

HS-PS1-3

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Bond Dissociation Energy and Activation Energy

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Bond Energies and Bond Lengths

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Coupled Reactions

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Energy Diagrams, Transition States, and Intermediates

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Types of Chemical Reactions: Anabolic and Catabolic

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Activation Energy

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Factors Influencing the Rate of Chemical Reactions

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Develop a model to illustrate that the release or absorption of energy from a chemical reaction system depends upon the changes in total bond energy.  Clarification Statement: Emphasis is on the idea that a chemical reaction is a system that affects the energy change.  Examples of models could include molecular-level drawings and diagrams of reactions, graphs showing the relative energies of reactants and products, and representations showing energy is conserved.

HS-PS1-4

How Electrons Fill Atomic Orbitals

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