Name: bond order and molecular orbitals
Uploaded: 2023-04-30T18:28:31.000Z
Duration: 2 min 40 s
Channel: JoVE

K12 Standards

Minnesota Academic Standards in Science

9-12

Chemistry

4.2.1 9C.3.2.1.2

MO Theory and Covalent Bonding

Теория молекулярных орбит описывает распределение электронов по молекуле, а не локализацию их в конкретных связях между атомами.
Конструктивная интерференция между синфазными атомными орбиталями соответствует большей электронной плотности между положительно заряженными ядрами, что делает молекулу более стабильной. Эта связывающая молекулярная орбиталь имеет меньшую энергию, чем любая из исходных атомных орбиталей.
Деструктивная интерференция между противофазными атомными орбиталями соответствует более низкой электронной плотности в узловой плоскости между ядрами, что делает молекулу менее стабильной. Эта антисвязывающая молекулярная орбиталь имеет более высокую энергию, чем атомные орбитали, и отмечена звездой или звездочкой.
Молекулярные орбитали классифицируются по тому, как атомные орбитали перекрываются. Лобовое сочетание атомных орбиталей вдоль межъядерной оси, такое как перекрытие между двумя s-орбиталями или двумя сквозными p-орбиталями, приводит к образованию сигма-молекулярных орбиталей. Сигма-орбитальная электронная плотность сосредоточена вокруг межъядерной оси.
Боковое перекрытие, такое как боковое перекрытие двух p-орбиталей, приводит к образованию пи-молекулярных орбиталей. Здесь электронная плотность сосредоточена на противоположных сторонах межъядерной оси.
Ориентация трех различных p-орбиталей означает, что обычно одна пара перекрывается встык, а две другие пары перекрываются вбок. Пи-связывающие орбитали обычно равны по энергии или вырождены, как и пи-антисвязывающие орбитали.
Теория молекулярных орбиталей предсказывает стабильность ковалентных связей из порядка связей молекулы, который представляет собой число электронов на связывающих орбиталях за вычетом числа электронов на антисвязывающих орбиталях, деленное на два.
Порядок облигаций больше нуля указывает на то, что может существовать одна или несколько ковалентных связей, в то время как нулевой порядок облигаций означает, что облигации не должны существовать.
Теория молекулярных орбит также полезна для многоатомных молекул, таких как бензол. Модель бензола Льюиса не может точно представить его делокализованные электроны, в то время как молекулярная орбитальная теория относит эти электроны к трем пи-связывающим молекулярным орбиталям, охватывающим все углеродное кольцо.

Этот текст адаптирован из <a href="https://openstax.org/books/chemistry-2e/pages/8-4-molecular-orbital-theory">Openstax, Химия 2e, раздел 8.4 Теория молекулярных орбит.</a>

Теория МО и ионно-дипольные силы

Теория молекулярных орбиталей описывает распределение электронов в молекулах аналогично распределению электронов по атомным орбиталям. Область пространства, в которой может находиться валентный электрон в молекуле, называется молекулярной орбиталью. Математически линейная комбинация пятиатомных орбиталей (ЛКАО) порождает молекулярные орбитали. Комбинации синфазных пятиатомных орбитальных волновых функций приводят к образованию областей с высокой вероятностью электронной плотности, тогда как противофазные волны создают узлы или области без электронной плотности.

Синфазная комбинация двух пятиатомных s-орбиталей на соседних атомах создает связывающую молекулярную орбиталь σs с более низкой энергией, в которой большая часть электронной плотности находится непосредственно между ядрами. Противофазное присоединение приводит к образованию разрыхляющей молекулярной орбитали σs* с более высокой энергией, на которой между ядрами имеется узел.

Аналогично, волновая функция p-орбиталей порождает два лепестка с противоположными фазами. Когда p-орбитали перекрываются на концах, они создают орбитали σ и σ*. Перекрытие “бок о бок” двух p-орбиталей приводит к образованию π-связывающих и π*-разрыхляющих молекулярных орбиталей.

Заполненная диаграмма молекулярных орбиталей показывает количество электронов на связывающих и разрыхляющих молекулярных орбиталях. Электрон участвует в связывающем взаимодействии только в том случае, если он занимает связывающую орбиталь. Общий вклад электронов в прочность связи молекулы определяется по порядку связи, который рассчитывается следующим образом:

<img alt="Figure1" src="/files/ftp_upload/11687/11687%20PT%20equation%201.png" style="text-align: center; height: 40px;">

Порядок связи является показателем прочности ковалентной связи; связь между двумя данными атомами становится прочнее с увеличением порядка связи. Если распределение электронов по молекулярным орбиталям дает нулевой порядок связи, стабильная связь не образуется.

Теория молекулярных орбиталей также полезна для многопятиатомных молекул. Модель Льюиса бензола (C6H6), который имеет плоскую гексагональную структуру с sp2-гибридизированными атомами углерода, не может точно представить его делокализованные электроны. Однако теория молекулярных орбиталей относит эти электроны к трем π-связывающим молекулярным орбиталям, охватывающим все углеродное кольцо. В результате образуется полностью занятой (6 электронов) набор связывающих молекулярных орбиталей, которые придают бензольному кольцу дополнительную термодинамическую и химическую стабильность.

Теория молекулярных орбиталей описывает распределение электронов в молекулах подобно распределению электронов на атомных орбиталях. ...

Education

JoVE Core

Organic Chemistry

Органическая химия

Unit 1

Раздел 1

Covalent Bonding and Structure

Ковалентная связь и структура

KEY TERMS AND CONCEPTS

ОСНОВНЫЕ ТЕРМИНЫ И ПОНЯТИЯ

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 electron density.

Two adjacent s orbitals can overlap to form a lower-energy sigma s bonding molecular orbital. In this orbital, most of the electron density lies between the nuclei. The out-of-phase combination forms a higher-energy sigma s* antibonding molecular orbital, which has a node between the nuclei.

p orbitals also build molecular orbitals. Their wave functions have two lobes with opposite phases. End-to-end overlap of p orbitals produces sigma and sigma* orbitals, while side-by-side overlap produces pi bonding and pi* antibonding molecular orbitals.

A filled molecular orbital diagram shows how many electrons occupy bonding and antibonding orbitals. An electron adds to bond strength only when it is in a bonding orbital. The overall effect is measured by bond order, which helps show how strong a covalent bond is.

As bond order increases, the bond between two atoms becomes stronger. If the molecular orbital arrangement gives a bond order of zero, a stable bond does not form. Molecular orbital theory also helps explain polyatomic molecules. In benzene, the Lewis model cannot show the delocalized electrons in its planar hexagonal ring of sp2 hybridized carbon atoms. Molecular orbital theory places those electrons in three pi bonding molecular orbitals spread across the whole ring. With 6 electrons filling these bonding orbitals, benzene gains extra thermodynamic and chemical stability.

Bond Order and Molecular Orbitals

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9C.3.2.1.3

4 Communicating reasons, arguments and ideas to others
 4.2 Obtaining, evaluating and communicating information
  4.2.1 Students will be able to read and interpret multiple sources to obtain information, evaluate the merit and validity of claims and design solutions, and communicate information, ideas, and evidence in a variety of formats.

Bond Order and Molecular Orbitals

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