Name: drawing lewis structures with valence electron counts
Uploaded: 2020-09-03T13:38:28.000Z
Duration: 2 min 54 s
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Lewis Structures of Molecular Compounds and Polyatomic Ions

Структуры Льюиса это упрощенные представления химических связей между атомами. Написание структур Льюиса для нейтральных соединений, таких как этен или аммиак, или для многоатомных ионов, таких как гидроксоний или фосфат, включает следующую последовательность шагов. Сначала рассчитаем общее количество валентных электронов в молекуле.Рассмотрим, например, этен. Каждый атом углерода имеет 4 валентных электрона, а каждый атом водорода имеет 1 валентный электрон;с 2 атомами углерода и 4 атомами водорода у этена всего 12 валентных электронов. Затем нарисуйте скелетную структуру молекулы.Обычно первый атом в химической формуле или менее электроотрицательный атом размещается по центру. Этот атом также называют центральным атомом, и он должен образовывать как минимум две связи. Атомы водорода или несколько электроотрицательных атомов расположены на концах вокруг центрального атома.Здесь атомы углерода занимают центральные позиции с окружающими их атомами водорода. Следующий шаг завершение октета. Во-первых, поместите пару электронов между каждыми двумя атомами, чтобы представить связывающие пары.Остальные электроны распределяются в виде неподеленных пар между концевыми и центральными атомами, чтобы удовлетворить конфигурациям дуэта или октета. В этом случае, поскольку водород достиг дуэта, оставшиеся электроны передаются атомам углерода. Теперь используйте неподеленные пары, чтобы сформировать множественные связи между атомами, у которых отсутствует октет.В этене атомы углерода еще не достигли октета. Следовательно, неподеленная пара перемещается в область связи, чтобы сформировать двойную связь, позволяющую всем атомам достичь стабильной электронной конфигурации. Чтобы записать структуру Льюиса для заряженных многоатомных ионов, таких как гидроксоний или фосфат, используется тот же подход, но с модификацией, учитывающей ионный заряд.Чтобы нарисовать структуру Льюиса для гидроксония, сначала вычислите количество валентных электронов на ионе. Поскольку ион гидроксония имеет положительный заряд, удалите один электрон из общего количества валентных электронов, уменьшив его до 8. Затем изобразите молекулу ее скелетной структурой.Теперь поместите электронную пару между каждым атомом, а затем неподеленную пару электронов на центральном атоме. Таким образом, атомы водорода и кислорода удовлетворяют дуэту и октету соответственно. Наконец, запишите в скобках структуру Льюиса с положительным зарядом в правом верхнем углу.

Льюисовские структуры молекулярных соединений и многоатомных ионов

Чтобы изобразить структуры Льюиса для сложных молекул и молекулярных ионов, рекомендуется выполнить пошаговую процедуру, как описано ниже: 
 
 <ol>
 <li>Определите общее количество электронов валентности (внешней оболочки). Для катионов вычтите по одному электрона за каждый положительный заряд. Для анионов добавьте по одному электрона каждый отрицательный заряд. </li>
 <li>Нарисуйте скелетную структуру молекулы или иона, расположив атомы вокруг центрального атома. (Как правило, наименее электронный элемент должен быть расположен в центре.) Соедините каждый атом с центральным атомом с помощью одной связи (одна электронная пара). </li>
 <li>Распределите оставшиеся электроны как одинокие пары на атомах терминала (кроме водорода), завершая октет вокруг каждого атома. </li>
 <li>Поместите все оставшиеся электроны на центральный атом. </li>
 <li>Переставьте электроны внешних атомов, чтобы сделать несколько связей с центральным атомом, чтобы получить октеты, где это возможно. </li>
 </ol>
 
 Например, в качестве примеров можно привести SiH4, CHO2–, NO+ и OF2, для которых можно применять эти общие рекомендации для определения структур Льюиса. 
 
 <ol>
 <li>Определите общее количество электронов валентности (внешней оболочки) в молекуле или ионе.  </li>
 </ol>
 
 Для молекулы типа SiH4 добавляется количество валентных электронов на каждом атоме в молекуле:  
 
 = [4 валентных e-/Si атом × 1 Si атом ] + [ 1 валентный e-/H атом × 4 H атомов] = 8 валентный e- 
 
 <h4>CHO2– </h4>
 
 Для отрицательного иона, например CHO2–, число валентных электронов на атомах добавляется к количеству отрицательных зарядов на ионе (один электрон получают за каждый единичный отрицательный заряд): 
 
 = [4 валентных e-/C atom × 1 C atom] + [1валентный e-/H atom × 1 H atom] + [6 валентных e-/O atom × 2 O atoms] + [1 дополнительный e-] = 18 валентных e- 
 
 <h4>NO+ </h4>
 
 Для положительного иона, например NO+, добавляются числа валентных электронов на атомах в ионе, за которыми следует вычитание количества положительных зарядов на ионе (один электрон теряется за каждый положительный заряд) из общего числа валентных электронов: 
 
 = [5 валентных e-/N atom × 1 N атом] + [6 валентных e-/O атом × 1 O атом] + [-1 e-] = 10 валентных e- 
 
 <h4>О 2 </h4>
 
 Поскольку О2 является нейтральной молекулой, количество валентных электронов просто добавляется: 
 
 = [6 валентных e-/O атом × 1 O атом] + [7 валентных e-/F атом × 2 F атома] = 20 валентных e- 
 
 <ol start=""2"">
 <li>Нарисуйте скелетную структуру молекулы или иона, располагая атомы вокруг центрального атома и соединяя каждый атом с центральным атомом с помощью одной (одной пары электронов) связи. (Обратите внимание, что ионы обозначены скобками вокруг конструкции, а ионный заряд за пределами скобок:) 
 <img alt=""Figure1"" src=""https://www.jove.com/files/core/CDNSource/lessons/1063/1063Image_1.svg"" style=""height: 80px;"" /> 
 В тех случаях, когда возможно несколько способов аранжировки атомов, как и CHO2, для выбора среди них верной используются экспериментальные доказательства. В целом, менее электроотрицательные элементы, скорее всего, будут центральными атомами. В CHO2 меньше электроотрицательный атом углерода занимает центральное положение с окружающими его атомами кислорода и водорода. Другие примеры включают P в POCl3, S в SO2 и Cl в ClO4.  Исключением является то, что водород почти никогда не является центральным атомом. Поскольку фтор является самым электроотрицательным элементом, он также не может быть центральным атомом. </li>
 <li>Распределите оставшиеся электроны как одинокие пары на атомах терминала (за исключением водорода), чтобы завершить валентные оболочки октетом электронов. (При отсутствии оставшихся электронов на SiH4 его структура остается неизменной.) 
 <img alt=""Figure2"" src=""https://www.jove.com/files/core/CDNSource/lessons/1063/1063image_2.svg"" style=""height: 80px;"" /></li>
 <li>Поместите все оставшиеся электроны на центральный атом.  
 <ul>
 <li>Для SiH4, CHO2 и NО+ не осталось электронов. Для 2 из 16 оставшихся электронов 12 размещаются, таким образом оставляя 4 электрона для размещения на центральном атоме: 
 <img alt=""Figure3"" src=""https://www.jove.com/files/core/CDNSource/lessons/1063/1063Image_3.svg"" style=""height: 40px;"" /></li>
 </ul>
 </li>
 <li>Переставьте электроны внешних атомов, чтобы сделать несколько связей с центральным атомом и получить октеты, где это возможно.  
 <ul>
 <li>SiH4: SI уже имеет октет, поэтому ничего не нужно делать.   
 <img alt="Figure4" src="https://www.jove.com/files/core/CDNSource/lessons/1063/1063Image_4.svg" style="height: 70px;" /></li>
 <li>CHO2-: Валентные электроны распределяются как неподелённые пары на атомах кислорода, но атом углерода не имеет октета.   
 <img alt="Figure5" src="https://www.jove.com/files/core/CDNSource/lessons/1063/1063Image_5.svg" style="height: 80px;" /></li>
 <li>Таким образом, одна неподелённая пара электронов жертвуется одним из кислородов атому углерода, образуя двойную связь. В зависимости от того, какой атом кислорода пожертвовал электроны, могут быть две возможные структуры, иначе называемые резонансные структуры.  </li>
 <li>NО+: Для этого иона добавляется восемь валентных электронов, но ни один из атомов не имеет октета. Невозможно добавить дополнительные электроны, так как общее количество электронов уже используется. В этом сценарии электроны должны быть перемещены, чтобы сформировать несколько связей. Атом азота имеет две неподелённые пары электронов, а атом кислорода имеет одну.   
 <img alt="Figure6" src="https://www.jove.com/files/core/CDNSource/lessons/1063/1063Image_6.svg" style="height: 55px;" /></li>
 <li>При этом октет по-прежнему не генерируют, поэтому необходимо переместить другую пару, чтобы сформировать тройную связь. 
  <img alt="Figure7" src="https://www.jove.com/files/core/CDNSource/lessons/1063/1063Image_7.svg" style="height: 50px;" /></li>
 <li>В О2 ничего не меняется, так как каждый атом уже имеет октет. 
 <img alt="Figure8" src="https://www.jove.com/files/core/CDNSource/lessons/1063/1063Image_8.svg" style="height: 40px;" /></li>
 </ul>
 </li>
 </ol>
 
 Этот текст адаптирован из <a href="https://openstax.org/books/chemistry-2e/pages/7-3-lewis-symbols-and-structures">Openstax, Химия 2е изд., раздел 7.3:Символы и структуры Льюиса.</a>

Чтобы нарисовать структуры Льюиса для сложных молекул и молекулярных ионов, полезно следовать пошаговой процедуре, как показано ниже:

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drawing lewis structures with valence electron counts

Drawing Lewis structures starts with counting valence electrons and building the atom arrangement step by step. Valence electrons are the outer-shell electrons that help show how atoms bond. For ions, a positive charge means one electron is removed for each plus sign, and a negative charge means one electron is added for each minus sign.

The next step is to sketch the skeleton structure. The least electronegative atom usually goes in the center, while hydrogen is almost never central. Fluorine also cannot be a central atom because it is the most electronegative element. Ions are written in brackets, with the charge shown outside the brackets.

After the skeleton is drawn, connect the atoms to the central atom with single bonds. Each single bond represents one electron pair. Then place the remaining electrons as lone pairs on the terminal atoms, except hydrogen, until those atoms have octets when possible. Any electrons still left go on the central atom.

If the central atom still does not have an octet, move lone pairs to make double or triple bonds. This can create resonance structures, which are different valid arrangements of electrons. For the formate ion, CHO2−, one oxygen can donate a lone pair to make a double bond with carbon. That gives two possible resonance forms.

The examples in the model show how the procedure works in real cases. SiH4 has 8 valence electrons total, so its structure does not need more changes once the single bonds are drawn. CHO2− has 18 valence electrons. NO+ has 10 valence electrons, and its atoms must be rearranged into a triple bond to satisfy octets as much as possible. OF2 has 20 valence electrons, and after the lone pairs are placed, its structure stays unchanged because each atom already has an octet.

Drawing Lewis Structures with Valence Electron Counts

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Chemical bonding explains how atoms and ions join to form molecules and compounds. American chemist Gilbert N. Lewis pioneered chemical bonding theory with the Lewis model. This model helps explain why bonds form, why some are stronger than others, and why compounds have fixed formulas like H2O instead of H4O. Ionic bonding involves ions, which are atoms or molecules with an electrical charge. A cation is a positive ion formed when a neutral atom loses one or more electrons from its valence...

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Chemical bonds form when atoms or ions interact in ways that lower the potential energy of a molecule. Gilbert N. Lewis created a simple model to show bond formation using valence electrons, the electrons in the outer shell. In this model, electrons are shared or transferred to help atoms reach a stable electron configuration. The Lewis model helps describe two common bond types. An ionic bond forms when electrons move from a metal to a nonmetal. A covalent bond forms when two nonmetals share...

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Ionic Compounds and Crystal Lattices

Ionic compounds form when positive and negative ions attract each other. A cation is a positive ion, and it forms when an atom loses one or more valence electrons. An anion is a negative ion, and it forms when an atom gains one or more valence electrons. Compounds made of these ions are called ionic compounds, or salts, and the force between them is an ionic bond. The properties of ionic solids help show how strong this attraction is. They usually form rigid, brittle crystals and have high...

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Lattice energy shows how strongly ions hold an ionic solid together. It is the energy required to separate one mole of a solid into its gaseous ions. For sodium chloride, this process requires 769 kJ per mole. The opposite convention may also be used, where lattice energy is the energy released when gaseous ions form the solid. In either case, a larger magnitude means a more stable ionic compound. For ionic crystals, lattice energy follows a pattern based on charge and ion size. The...

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Shared Electrons in Covalent Molecules

Covalent molecules form when atoms share electrons instead of transferring them. This bonding usually happens between nonmetal atoms with similar tendencies to attract electrons. Because the atoms pull on the electrons in a shared way, the bond is held together by mutual attraction. Covalent compounds often have physical properties that differ from ionic compounds. The attraction between neutral molecules is weaker than the attraction between charged ions. As a result, many covalent compounds...

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How Bond Polarity Shapes Molecular Dipoles

Bond polarity depends on electronegativity, the ability of an atom to attract shared electrons. The absolute value of the electronegativity difference, or ΔEN, gives a rough guide to bond type. When the difference is very small or zero, the bond is covalent and nonpolar. When it is larger, the bond may be polar covalent or ionic. Examples show how this works in real compounds. H–H has a ΔEN of 0, so it is nonpolar. H–Cl has a ΔEN of 0.9, so it is polar covalent. Na–Cl has a ΔEN of 2.1, so it...

lewis structures of molecular compounds and polyatomic ions

Drawing Lewis structures starts with counting valence electrons and building the atom arrangement step by step. Valence electrons are the outer-shell electrons that help show how atoms bond. For ions, a positive charge means one electron is removed for each plus sign, and a negative charge means one electron is added for each minus sign. The next step is to sketch the skeleton structure. The least electronegative atom usually goes in the center, while hydrogen is almost never central. Fluorine...

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Resonance Forms in Lewis Structures

Resonance forms in Lewis structures help explain molecules and ions that cannot be shown correctly with just one drawing. In nitrite ion, NO2−, the Lewis structure can be drawn in two ways. The N-O and N=O bonds switch positions in the two drawings, but the atoms stay in the same arrangement. If nitrite really had one single bond and one double bond, the two bond lengths should be different. A double bond is shorter and stronger than a single bond between the same atoms. Experiments show that...

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Lewis Structures with Formal Charge

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Electron Sea Model in Metals

Electron sea model explains how metallic bonding holds metals together. In this model, most metal atoms do not have enough valence electrons to form ionic or covalent bonds in the usual way. Their valence electrons are held loosely because metals have low electronegativity, and their ionization energy is low, so electrons are removed easily. When a metal atom loses these outer electrons, it becomes a positively charged metal ion. The freed electrons form a delocalized electron cloud around...

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Learn the step-by-step method for drawing Lewis structures, counting valence electrons, placing lone pairs, and handling ions, resonance, and multiple bonds.

Drawing Lewis Structures with Valence Electron Counts | JoVE K12

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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 ...

The Aufbau Principle and Hund's Rule

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

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.

PS.Chm.1.3

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

Atoms &mdash; and the protons, neutrons, and electrons that compose them &mdash; are extremely small. For example, a carbon atom weighs less than 2 ...

Molar Mass

The identity of a substance is defined not only by the types of atoms or ions it contains but by the quantity of each type of atom or ion. For example, ...

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

The elements in groups of the periodic table exhibit similar chemical behavior. This similarity occurs because the members of a group have the same number ...

Ionic Radii

Ionic radius is the measure used to describe the size of an ion. A cation always has fewer electrons and the same number of protons as the parent atom; it ...

Ionization Energy

The amount of energy required to remove the most loosely bound electron from a gaseous atom in its ground state is called its first ionization energy ...

Electron Affinity

The electron affinity (EA) is the energy change for adding an electron to a gaseous atom to form an anion (negative ion).



This process can be either ...

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

Types of Chemical Bonds

Chemical bonding theories were pioneered by American chemist Gilbert&#xA0;N. Lewis. He developed a model called the Lewis model to explain the type and ...

Lewis Symbols and the Octet Rule

Chemical bonds are complex interactions between two or more atoms or ions, which reduce the potential energy of the molecule. Gilbert N. Lewis developed a ...

Ionic Bonding and Electron Transfer

Ions are atoms or molecules bearing an electrical charge. A cation (a positive ion) forms when a neutral atom loses one or more electrons from its valence ...

The Born-Haber Cycle

Lattice Energy&#xA0;

An ionic compound is stable because of the electrostatic attraction between its positive and negative ions. The lattice energy of a ...

Trends in Lattice Energy: Ion Size and Charge

An ionic compound is stable because of the electrostatic attraction between its positive and negative ions. The lattice energy of a compound is a measure ...

Covalent Bonding and Lewis Structures

Compared to ionic bonds, which results from the transfer of electrons between metallic and nonmetallic atoms, covalent bonds result from the mutual ...

Electronegativity

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

Electronegativity values of ...

Bond Polarity, Dipole Moment, and Percent Ionic Character

Bond Polarity

The absolute value of the difference in electronegativity (&Delta;EN) of two bonded atoms provides a rough measure of the polarity to be ...

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

Analyze the structure of atoms and isotopes.

Drawing Lewis Structures with Valence Electron Counts

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