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

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High School

Chemistry

Matter And Its Interactions

HS-PS1-1

Electron Configurations

La configuration électronique d’un atome représente la distribution des électrons entre ses orbitales atomiques. Le principe d’exclusion de Pauli, la règle de multiplicité maximale de Hund et le principe d’Aufbau peuvent être étendus pour envisager la configuration électronique de n’importe quel élément.
Le principe d’Aufbau stipule que dans l’état fondamental, les orbitales atomiques remplissent un ordre croissant d’énergie. Les énergies relatives des orbitales atomiques sont rationalisées par les interactions de Coulomb, l’effet de blindage et la pénétration orbitale.
Considérons la configuration électronique du carbone, un élément de numéro atomique six et donc neutre avec six électrons. L’orbitale1 s, qui a l’énergie la plus faible, est remplie en premier.
Selon le principe d’exclusion de Pauli, deux électrons d’un atome ne peuvent pas avoir le même ensemble de quatre nombres quantiques. Comme les électrons d’une même orbitale ont les mêmes nombres quantiques principaux, azimutaux et magnétiques, ils doivent avoir des nombres quantiques de spin différents.
Comme le nombre quantique de spin n’a que deux valeurs possibles, une orbitale ne peut accueillir que deux électrons, de spins opposés.
Bien que l’énergie augmente avec le nombre de couches, la plus grande pénétration des orbitales s abaisse l’énergie des orbitales s par rapport à celle des orbitales p.
Par conséquent, les deux électrons suivants occupent l’orbitale 2s et le cinquième entre dans la sous-couche 2p. Comme les orbitales d’une sous-couche sont présumées dégénérées, le cinquième électron peut entrer dans l’une des trois orbitales 2p dégénérées.
Le sixième électron suit la règle de multiplication maximale de Hund et occupe individuellement une orbitale dégénérée plutôt qu’une appariement. Par conséquent, pour le carbone, les deux électrons 2p occupent deux orbitales différentes et ont des spins parallèles.
Le diagramme de configuration électronique révèle que le carbone a deux électrons, appelés électrons de cœur, dans la couche interne et quatre électrons, appelés électrons de valence, dans les couches les plus externes.
L’ordre des orbitales atomiques par rapport à leurs énergies peut être mémorisé à l’aide de tels diagrammes, où la trajectoire de la flèche révèle la séquence dans laquelle les électrons sont attribués aux orbitales.
Cependant, cette progression devient plus complexe à mesure que les orbitales d et f sont introduites et la séquence devient moins prévisible pour les éléments de transition, les lanthanides et les actinides.

configurations électroniques

Les configurations électroniques et les diagrammes orbitaux peuvent être déterminés en appliquant le principe d'Aufbau (chaque électron ajouté occupe la sous-couche de la plus basse énergie disponible), le principe d'exclusion de Pauli (deux électrons ne peuvent pas avoir le même ensemble de quatre nombres quantiques) et la règle de Hund de multiplicité maximale. (Dans la mesure du possible, les électrons conservent des spins non appariés dans des orbitales dégénérées.)

Les énergies relatives des sous-couches déterminent l'ordre dans lequel les orbitales atomiques sont remplies (1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, la tendance du pouvoir de pénétration d’un électron peut être représentée de la manière suivante :

1s > 2s > 2p > 3s > 3p > 4s > 3d > 4p > 5s > 4d > 5p > 6s > 4f....

L'effet de blindage et de pénétration orbitale est important et un électron 4s peut avoir une énergie inférieure à celle d'un électron 3d.

Les électrons situés sur les orbitales les plus externes, appelés électrons de valence, sont responsables de la majeure partie du comportement chimique des éléments. Dans le tableau périodique, les éléments ayant des configurations électroniques de valence analogues apparaissent généralement dans le même groupe.

Il existe quelques exceptions à l'ordre de remplissage prévu, en particulier lorsque des orbitales à moitié ou complètement remplies peuvent être formées. Dans le cas du Cr et du Cu, les sous-couches à moitié remplies et complètement remplies représentent apparemment des conditions de stabilité privilégiées. Cette stabilité est telle que l'électron se déplace de l'orbitale 4s vers l'orbitale 3d pour obtenir la stabilité supplémentaire d'une sous-couche 3d à moitié remplie (en Cr) ou d'une sous-couche 3D remplie (en Cu). D'autres exceptions se produisent également. Par exemple, le niobium (Nb, numéro atomique 41) devrait avoir la configuration électronique [Kr]5s24d3. Cependant, expérimentalement, sa configuration électronique à l’état fondamental est en réalité [Kr]5s14d4. Nous pouvons rationaliser cette observation en disant que les répulsions électron-électron ressenties par l’appariement des électrons dans l’orbitale 5s sont plus grandes que l’écart d’énergie entre les orbitales 5s et 4d.

<a href="https://openstax.org/books/chemistry-2e/pages/6-4-electronic-structure-of-atoms-electron-configurations">Ce texte est adapté de Openstax, Chemistry 2e, Section 6.4 : Structure électronique des atomes.</a>

Configurations électroniques

Les configurations électroniques et les diagrammes orbitaux peuvent être déterminés en appliquant le principe d’Aufbau (chaque électron ajouté occupe la sous-couche de ...

Education

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Organic Chemistry

Chimie organique

Unit 1

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Covalent Bonding and Structure

Liaison covalente et structure

KEY TERMS AND CONCEPTS

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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?

Carbon Compounds in Everyday Life

Organic chemistry focuses on carbon compounds called organic compounds. These compounds may come from living things or be made by chemists. Carbon is the main element in them, and it bonds with other carbon atoms as well as hydrogen, oxygen, nitrogen, and sulfur. Carbon can form up to four strong bonds. That ability lets carbon atoms build chains and rings of many sizes, shapes, and levels of complexity. Not every carbon-containing substance is classified as organic. Carbonates, cyanides, and...

electronic structure of atoms

Atomic Orbitals and Quantum Numbers

Atomic orbitals and quantum numbers describe how electrons are arranged around the nucleus. An atom has a small, positively charged nucleus with electrons surrounding it. The nucleus contains protons and neutrons, and the quantum mechanical model gives a better picture of this structure than older models. Protons carry a +1 charge and have a mass of 1.0073 amu. Neutrons have no charge and have a mass of 1.0087 amu. Electrons carry a −1 charge and have a mass of 0.00055 amu. The number of...

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...

polar covalent bonds

Electronegativity in Polar Covalent Bonds

Polar covalent bonds form when atoms share electrons unevenly. This happens in covalent bonds between atoms that have similar, but not identical, tendencies to attract electrons. Nonmetal atoms often make these bonds with other nonmetals. Hydrogen gas, H2, is a simple example of a covalent bond between two hydrogen atoms. A covalent bond begins when two atoms move close enough for their valence orbitals to overlap. In hydrogen, the 1s orbitals overlap and the single electrons on each atom...

lewis structures and formal charges

Lewis Structures, Bonds, and Formal Charge

Lewis structures show how atoms share electrons in covalent bonds and in polyatomic ions. Lewis symbols help draw these structures, and the periodic table helps predict how many valence electrons an atom has and how many bonds it can form to reach an octet. Group 18 elements, such as argon and helium, already have filled electron configurations. They rarely take part in chemical bonding. Atoms in group 17, such as bromine and iodine, need one more electron to reach an octet, so they can form...

vsepr theory

Molecular Shapes from Electron Pairs

Molecular shape can be predicted from the way electron pairs arrange around a central atom. Valence shell electron-pair repulsion theory, or VSEPR theory, uses a Lewis structure to count bonding pairs and lone electron pairs. The main idea is simple: electron pairs spread out as far as possible to reduce repulsion. In a molecule, the valence electrons around the central atom can be bonding pairs or lone pairs. Bonding pairs are found mainly between bonded atoms. Lone pairs stay on one atom and...

molecular geometry and dipole moments

Predicting Shape and Molecular Polarity

Molecular shape and molecular polarity can be predicted with VSEPR theory and dipole moments. Start by writing the Lewis structure of the molecule or polyatomic ion. Then count the electron groups around the central atom. Lone pairs and bonds both count as electron groups, and a single, double, or triple bond each counts as one region of electron density. Use that count to identify the electron-pair geometry. Next, use the number of lone pairs to determine the molecular structure. If more than...

resonance and hybrid structures

Resonance in Nitrite and Carbonate Ions

Resonance explains why some Lewis structures do not match the measured properties of a molecule or ion. When two or more Lewis structures have the same atom arrangement, the real electron distribution is an average of the structures shown. The nitrite anion, NO2−, is one example. It can be drawn in two ways, with the N–O and N=O bonds in different places. If nitrite had one single bond and one double bond, the two bond lengths should be different, because a double bond is shorter and stronger...

valence bond theory and hybridized orbitals

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Covalent Bond Overlap and Orbital Hybridization

Covalent bonds form when orbitals on two atoms overlap and their single electrons pair up. The bond is stronger when the overlap is greater. Maximum overlap happens when the orbitals meet along a direct line between the two nuclei. A sigma bond, or σ bond, is the kind of covalent bond found in a single bond in a Lewis structure. In a σ bond, electron density is concentrated along the internuclear axis. A pi bond, or π bond, forms when two p orbitals overlap side by side. In a π bond, the...

mo theory and covalent bonding

Bond Order and Molecular Orbitals

Molecular orbital theory explains how electrons are arranged in molecules. It treats electron regions in a molecule much like atomic orbitals in an atom. A molecular orbital is the region of space where a valence electron is most likely to be found. Molecular orbitals form through the linear combination of atomic orbitals, or LCAO. When atomic wave functions combine in phase, they create regions with high electron density. When they combine out of phase, they create nodes, or regions with no...

intermolecular forces and physical properties

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How Intermolecular Forces Affect Matter

Intermolecular forces are the attractions between molecules, and they help explain how matter behaves. These forces influence melting points, boiling points, and whether substances mix with one another. Water has a much higher boiling point than hexane, which shows that water has stronger intermolecular forces. Three main kinds of intermolecular interactions are covered here. They include ion-dipole forces, dipole-dipole interactions, and van der Waals forces, which include London dispersion...

solubility

Why Some Substances Dissolve in Water

Solution formation depends on how solutes and solvents interact. A solution is a homogeneous mixture made of a solvent, which is the larger part, and a solute, which is the smaller part. The physical state of a solution is usually the same as the solvent. Concentration is often described as dilute for a low amount of solute or concentrated for a high amount of solute. When a solute dissolves, its particles are surrounded by solvent species. This process is called solvation. If water is the...

introduction to functional groups

Functional Groups and Organic Molecules

Functional groups are groups of atoms with specific chemical properties that appear within organic molecules. They are sometimes shown as "R." These groups sit along the carbon backbone of macromolecules, which can form chains or rings of carbon atoms. Functional groups help functionalize a compound. That means they can give the molecule different physical and chemical properties. In organic chemistry, this change in group structure can strongly affect how a compound behaves. Common...

overview of advanced functional groups

Advanced Functional Groups in Organic Chemistry

Advanced functional groups are atom clusters in organic molecules that give those molecules specific chemical properties. They are sometimes shown as "R" in formulas. By changing the atoms attached to a carbon skeleton, a functional group can make a compound behave differently in physical and chemical reactions. This topic covers several major classes of organic compounds with advanced functional groups. Aromatic hydrocarbons include benzene. Phenol is another example and is also called...

key terms and concepts

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

organic chemistry

JoVE Core Chimie Organique explique les concepts de base de la stéréochimie, des réactions et des mécanismes à l'aide de leçons vidéo animées concises et faciles à comprendre. De plus, les vidéos de scientifiques à l'action présentent l'application de ces concepts grâce à des expériences de recherche classiques et originales réalisées dans des laboratoires modernes.

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Matter and Its Interactions

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Periodic Classification of the Elements

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the-periodic-table-and-organismal-elements

The Periodic Table And Organismal Elements

Atomic Structure

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Electronic Structure of Atoms


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Electron Configuration of Multielectron Atoms

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Atomic Radii and Effective Nuclear Charge

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The Aufbau Principle and Hund's Rule

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Coulomb's Law

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Sources and Properties of Electric Charge

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Electric Fields

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Intermolecular Forces

Atoms and molecules interact through bonds (or forces): intramolecular and intermolecular. The forces are electrostatic as they arise from interactions ...

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

In ordinary chemical reactions, the nucleus &mdash; which contains the protons and neutrons of each atom and thus identifies the element &mdash; remains ...

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

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

Chemical Reactions

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Chemical Equations

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Introduction to Chemical Reactions

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Temperature Dependence on Reaction Rate

The Collision Theory
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Stoichiometry, Product Yield, and Limiting Reactants

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Physical and Chemical Properties of Matter

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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).

Intermolecular vs Intramolecular Forces

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Comparison Between Electrical And Gravitational Forces

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Intermolecular Forces and Physical Properties



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Electric Field of Two Equal and Opposite Charges

Atoms generally contain the same number of positively and negatively charged particles, protons, and electrons. Hence, they are electrically neutral. ...

Electric Field

Consider two point charges, each exerting Coulomb force on the other. It is possible to describe the Coulomb interaction via an intermediate step by ...

Ionic Bonding and Electron Transfer

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Electric Dipoles and Dipole Moment

Consider two charges of equal magnitude but opposite signs. If they cannot be separated by an external electric field, the system is called a permanent ...

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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What is Metabolism?

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Types of Chemical Reactions: Anabolic and Catabolic

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Activation Energy

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Energy Transfer in 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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