Name: shared electrons in covalent molecules
Uploaded: 2020-09-03T13:38:28.000Z
Duration: 2 min 46 s
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Chemistry

Matter and its Interactions

PS.Chm.3.1

Covalent Bonding and Lewis Structures

Kovalente Bindungen zwischen Nichtmetallen werden durch den Austausch von Valenzelektronen gebildet. Aber warum tauschen Nichtmetalle lieber Elektronen aus, anstatt sie wie in ionischen Bindungen zu übertragen? Nichtmetalle haben hohe Ionisationsenergien, was es schwierig macht, Valenzelektronen von einem Atom auf ein anderes zu übertragen.Betrachten Sie ein Molekül Ammoniak. Das Stickstoffatom benötigt drei weitere Elektronen, um ein Oktett zu erreichen, und das Wasserstoffatom benötigt ein Elektron, um ein Duett zu erreichen. Daher verbindet sich das Stickstoffatom mit drei Wasserstoffatomen, so dass sowohl Stickstoff als auch Wasserstoff stabile Elektronenkonfigurationen erreichen.Da jede dieser Bindungen ein Elektronenpaar teilt, wird sie als Einfachbindung bezeichnet. Das gemeinsame Elektronenpaar in der kovalenten Bindung wird als Bindungspaar bezeichnet. Die Valenzelektronen, die nicht an der Bindung beteiligt sind, werden als freies Elektronenpaar oder nicht-bindende Elektronen bezeichnet.Mit 6 Valenzelektronen benötigen Sauerstoffatome zwei weitere Elektronen, um ein Oktett zu erreichen. Daher teilen sich zwei Sauerstoffatome Zwei-Elektronen-Paare und bilden eine Doppelbindung. Stickstoff hingegen teilt sich die drei ungepaarten Elektronen in seiner zweiatomigen Form, wodurch eine Dreifachbindung entsteht.Einfach-und Mehrfachbindungen unterscheiden sich in Bezug auf Bindungslänge und stärke. Dreifachbindungen sind kürzer als Doppelbindungen, welche wiederum kürzer sind als Einfachbindungen. Die Bindungsstärke nimmt mit der Mehrfachbindung zu.Aus diesem Grund ist es schwierig, die Dreifachbindung in Stickstoff zu brechen, wodurch sie relativ reaktionsunfähig wird. Das Lewis-Modell hilft bei der Vorhersage der Struktur und Stabilität von Molekülen. Nach dem Lewis-Modell ist H2O ein stabiles Molekül, da sowohl Sauerstoff als auch Wasserstoff eine stabile Elektronenkonfiguration erreicht haben.Wenn der Sauerstoff Elektronen mit nur einem Wasserstoffatom teilt, ist das resultierende OH-Molekül nicht stabil, da Sauerstoff nur 7 Valenzelektronen hat und das Oktett nicht erreichen kann. Wenn dem Sauerstoff jedoch ein zusätzliches Elektron hinzugefügt wird, um das Oktett zu vervollständigen, wird das resultierende Hydroxid-Ion mit einer negativen Ladung stabil. Kovalente Bindungen sind gerichtet, weil die gemeinsamen Elektronen zwei bestimmte Atompaare verbinden.Im Gegensatz dazu sind ionische Bindungen nicht gerichtet und halten mehrere Ionen im Gitter. Daher werden in einer kovalenten Verbindung die einzelnen Moleküle als fundamentale Einheiten betrachtet.

kovalente bindung und lewis-strukturen

Im Vergleich zu ionischen Bindungen, die durch die Übertragung von Elektronen zwischen metallischen und nichtmetallischen Atomen entstehen, entstehen kovalente Bindungen durch die gleichzeitige Anziehung der Atomen durch ein „gemeinsames“ Elektronenpaar. Kovalente Bindungen werden zwischen zwei Atomen gebildet, wenn beide ähnliche Tendenzen Elektronen anzuziehen aufweisen (d. h. wenn beide Atome identische oder ziemlich ähnliche Ionisierungsenergien und Elektronenaffinitäten haben). <h4>Physikalische Eigenschaften kovalenter Verbindungen</h4> Verbindungen, die kovalente Bindungen enthalten, weisen andere physikalische Eigenschaften auf als ionische Verbindungen. Da die Anziehung zwischen elektrisch neutralen Atomen schwächer ist als die zwischen elektrisch geladenen Ionen, haben kovalente Verbindungen im Allgemeinen viel niedrigere Schmelz- und Siedepunkte als ionische Verbindungen. Tatsächlich sind viele kovalente Verbindungen bei Raumtemperatur flüssig oder gasförmig und in ihrem festen Zustand typischerweise viel weicher als ionische Feststoffe. Während ionische Verbindungen in Wasser gelöst gute Stromleiter sind, sind die meisten kovalenten Verbindungen in Wasser unlöslich und da sie elektrisch neutral sind, sind sie in jedem Zustand schlechte Stromleiter. <h4>Bildung kovalenter Bindungen</h4> Nichtmetallatome bilden häufig kovalente Bindungen mit anderen Nichtmetallatomen. Beispielsweise enthält das Wasserstoffmolekül H2 eine kovalente Bindung zwischen seinen beiden Wasserstoffatomen. Nähern sich zwei getrennte Wasserstoffatome mit einer bestimmten potentiellen Energie einander, beginnen sich ihre Valenzorbitale (1s) zu überlappen. Die einzelnen Elektronen jedes Wasserstoffatoms interagieren dann mit beiden Atomkernen und nehmen den Raum um beide Atome ein. Die starke Anziehungskraft jedes gemeinsamen Elektrons auf beide Kerne stabilisiert das System und die potentielle Energie nimmt mit abnehmendem Bindungsabstand ab. Nähern sich die Atome weiter einander an, beginnen sich die positiven Ladungen in den beiden Kernen gegenseitig abzustoßen und die potentielle Energie steigt. Die Bindungslänge wird durch den Abstand bestimmt, bei dem die niedrigste potentielle Energie erreicht wird. Es ist hierbei zu bedenken, dass zum Aufbrechen chemischer Bindungen Energie zugeführt werden muss (ein endothermer Prozess), während bei der Bildung chemischer Bindungen Energie freigesetzt wird (ein exothermer Prozess). Im Fall von H2 ist die kovalente Bindung sehr stark; Es muss eine große Energiemenge, 436 kJ, zugeführt werden, um die Bindungen in einem Mol Wasserstoffmoleküle aufzubrechen und die Atome zu trennen: <img alt="Eq1" src="https://www.jove.com/files/core/CDNSource/lessons/1059/1059_Equations-01.png" style="height: 40px;"> Umgekehrt wird die gleiche Energiemenge freigesetzt, wenn aus zwei Mol H-Atomen ein Mol H2-Moleküle entsteht: <img alt="Eq2" src="https://www.jove.com/files/core/CDNSource/lessons/1059/1059_Equations-02.png" style="height: 40px;"> <h4>Lewis-Strukturen</h4> Lewis-Strukturen und Symbole werden verwendet um die Bildung kovalenter Bindungen in Molekülen und mehratomigen Ionen zu beschreiben. Wenn beispielsweise zwei Chloratome ein Chlormolekül bilden, teilen sie sich ein Elektronenpaar: <img alt="Figure1" src="https://www.jove.com/files/core/CDNSource/lessons/1059/1059Image_1.svg" style="height: 90px;"> Die Lewis-Strukturformel deutet an, dass jedes Cl-Atom über drei Elektronenpaare verfügt, die nicht für die Bindung verwendet werden (die sogenannten freien Elektronenpaare) und über ein gemeinsames Elektronenpaar (das zwischen die Atomen gesetzt wird). Manchmal wird ein Strich (oder eine Linie) verwendet um ein gemeinsames Elektronenpaar anzuzeigen: Cl-Cl. <ul> <li>Ein einzelnes gemeinsames Elektronenpaar wird als Einfachbindung bezeichnet. Jedes Cl-Atom interagiert mit acht Valenzelektronen: den sechs in den freien Elektronenpaaren und den zwei in der Einfachbindung.</li> <li>Allerdings muss ein Atompaar möglicherweise mehr als ein Elektronenpaar teilen, um das erforderliche Oktett zu erreichen. Eine Doppelbindung entsteht, wenn zwei Elektronenpaare zwischen einem Atompaar geteilt werden, z. B. zwischen den Kohlenstoff- und Sauerstoffatomen in CH2O (Formaldehyd) und zwischen den beiden Kohlenstoffatomen in C2H4 (Ethylen). <img alt="Figure2" src="https://www.jove.com/files/core/CDNSource/lessons/1059/1059Image_2.svg" style="height: 120px;"></li> <li>Eine Dreifachbindung entsteht, wenn drei Elektronenpaare von einem Atompaar gemeinsam genutzt werden, wie bei Kohlenmonoxid (CO) und dem Cyanidion (CN–). <img alt="Figure3" src="https://www.jove.com/files/core/CDNSource/lessons/1059/1059Image_3.svg" style="height: 90px;"></li> </ul> Das Periodensystem kann verwendet werden, um die Anzahl der Valenzelektronen in einem Atom und die Anzahl der Bindungen vorherzusagen, die gebildet werden müssen, um ein Oktett zu erreichen. Elemente der Gruppe 18 wie Argon und Helium haben eine gefüllte Elektronenkonfiguration und nehmen daher selten an chemischen Bindungen teil. Atome der Gruppe 17 wie Brom oder Jod benötigen nur ein Elektron, um das Oktett zu erreichen. Daher können Atome der Gruppe 17 eine einzelne kovalente Bindung eingehen. Die Atome der Gruppe 16 benötigen 2 Elektronen, um ein Oktett zu erreichen; Daher können sie zwei kovalente Bindungen bilden. Ebenso benötigt Kohlenstoff, der zur Gruppe 14 gehört, vier weitere Elektronen, um ein Oktett zu erreichen; Somit kann Kohlenstoff vier kovalente Bindungen bilden. Dieser Text wurde adaptiert von <a href="https://openstax.org/books/chemistry-2e/pages/7-2-covalent-bonding">Openstax, Chemistry 2e, Section 7.2: Covalent Bonds</a> and <a href="https://openstax.org/books/chemistry-2e/pages/7-3-lewis-symbols-and-structures">Openstax, Chemistry 2e, Section 7.3: Lewis Symbols and Structures.</a>

Kovalente Bindung und Lewis-Strukturen

Im Vergleich zu ionischen Bindungen, die durch den Transfer von Elektronen zwischen metallischen und nichtmetallischen Atomen entstehen, ergeben sich kovalente Bindungen aus der gegenseitigen ...

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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 have lower melting and boiling points, and some are liquids or gases at room temperature. When they are solids, they are usually softer than ionic solids. Most covalent compounds are also insoluble in water and are poor conductors of electricity in any state.

Hydrogen gas shows how a covalent bond forms. In H2, two hydrogen atoms approach each other and their 1s valence orbitals overlap. Each electron interacts with both nuclei, which lowers the potential energy and stabilizes the molecule. If the atoms move too close, the positive nuclei repel each other, and the potential energy rises. The bond length is the distance where the potential energy is lowest. Forming a bond releases energy, so it is an exothermic process. Breaking a bond requires energy, so it is endothermic. For H2, 436 kJ must be added to break one mole of molecules, and the same amount is released when one mole forms from hydrogen atoms.

Lewis structures show covalent bonding with Lewis symbols. In chlorine gas, Cl2, the two chlorine atoms share one pair of electrons. Each chlorine atom also has three lone pairs, which are pairs not used in bonding. A shared pair can be shown with a dash, so Cl—Cl represents a single bond. Each chlorine atom then has eight valence electrons around it, including lone pairs and the bonding pair.

Some atoms need more than one shared pair to reach an octet, which is a full outer shell of eight electrons. A double bond forms when two pairs of electrons are shared, as in CH2O and C2H4. A triple bond forms when three pairs are shared, as in CO and CN−. The periodic table helps predict bonding patterns from valence electrons. Group 18 elements such as argon and helium already have filled electron configurations and rarely bond. Group 17 atoms, such as bromine and iodine, usually need one electron and can form one covalent bond. Group 16 atoms need two electrons and can form two covalent bonds. Carbon in group 14 needs four more electrons, so it can form four covalent bonds.

Shared Electrons in Covalent Molecules

types of chemical bonds

Chemical Bonding and Bond Properties

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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Lattice Energy Patterns in Ionic Solids

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...

covalent bonding and lewis structures

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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lewis structures of molecular compounds and polyatomic ions

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Explore how nonmetal atoms share electrons to form covalent bonds, from H2 and Cl2 to double and triple bonds, Lewis structures, and bonding patterns.

Shared Electrons in Covalent Molecules | JoVE K12

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Shared Electrons in Covalent Molecules

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