Name: shared electrons in covalent molecules
Uploaded: 2020-09-03T13:38:28.000Z
Duration: 2 min 46 s
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K12 Standards

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North Carolina K-12 Science Standard Course of Study

9-12

Chemistry

Matter and its Interactions

PS.Chm.3.1

Covalent Bonding and Lewis Structures

Los enlaces covalentes se forman entre no metales al compartir electrones de valencia. Pero, por qué los no metales prefieren compartir electrones en lugar de transferirlos como en los enlaces iónicos? Los no metales tienen altas energías de ionización que dificultan la transferencia de electrones de valencia de un átomo a otro.Considere una molécula de amoníaco. El átomo de nitrógeno requiere tres electrones más para alcanzar un octeto, y el átomo de hidrógeno necesita un electrón para alcanzar un dúo. Por lo tanto, el átomo de nitrógeno se une a tres átomos de hidrógeno de manera que tanto el nitrógeno como el hidrógeno alcanzan configuraciones electrónicas estables.Como cada uno de estos enlaces comparte un par de electrones, se llama enlace simple. El par de electrones compartido en en el enlace covalente se denomina par de enlace. Los electrones de valencia que no participan en el enlace se denominan par solitario o electrones no enlazados.Con 6 electrones de valencia, los átomos de oxígeno necesitan dos electrones más para alcanzar un octeto. Por lo tanto, dos átomos de oxígeno comparten pares de dos electrones formando un doble enlace. El nitrógeno, por otro lado, comparte los tres electrones desapareados en su forma diatómica, creando un triple enlace.Los enlaces simples y múltiples difieren en términos de longitud y resistencia del enlace. Los enlaces triples son más cortos que los enlaces dobles, que son más cortos que los enlaces simples. La fuerza de la unión aumenta con la multiplicidad de la unión.Por eso es difícil romper el triple enlace en el nitrógeno, lo que lo hace relativamente poco reactivo. El modelo de Lewis ayuda a predecir la estructura y estabilidad de las moléculas. Según el modelo de Lewis, el agua es una molécula estable;porque tanto el oxígeno como el hidrógeno han logrado configuraciones electrónicas estables.Si el oxígeno comparte electrones con un solo átomo de hidrógeno, la molécula de hidróxido resultante no es estable, ya que el oxígeno tiene solo 7 electrones de valencia y no puede alcanzar el octeto. Sin embargo, si se agrega un electrón extra al oxígeno para completar el octeto, el ion hidróxido resultante se vuelve estable con una carga negativa. Los enlaces covalentes son direccionales porque los electrones compartidos unen dos pares específicos de átomos.Por el contrario, los enlaces iónicos no son direccionales y contienen varios iones en la red. Por tanto, en un compuesto covalente, las moléculas individuales se consideran unidades fundamentales.

enlaces covalentes y estructuras de lewis

En comparación con los enlaces iónicos, que resultan de la transferencia de electrones entre átomos metálicos y no metálicos, los enlaces covalentes resultan de la atracción mutua de átomos para un par de electrones “compartido”. 
 
 Los enlaces covalentes se forman entre dos átomos cuando ambos tienen tendencias similares para atraer electrones hacia sí mismos (es decir, cuando ambos átomos tienen energías de ionización y afinidades electrónicas idénticas o bastante similares). 
 
 <h4>Propiedades físicas de los compuestos covalentes</h4>
 
 Los compuestos que contienen enlaces covalentes presentan propiedades físicas diferentes a aquellas de los compuestos iónicos. Debido a que la atracción entre moléculas, que son eléctricamente neutras, es más débil que la que existe entre los iones cargados eléctricamente, los compuestos covalentes generalmente tienen puntos de fusión y ebullición mucho más bajos que los compuestos iónicos. De hecho, muchos compuestos covalentes son líquidos o gases a temperatura ambiente y, en sus estados sólidos, son normalmente mucho más suaves que los sólidos iónicos. Además, mientras que los compuestos iónicos son buenos conductores de la electricidad cuando se disuelven en el agua, la mayoría de los compuestos covalentes son insolubles en agua; ya que son eléctricamente neutros, son malos conductores de la electricidad en cualquier estado.
 
 <h4>Formación de enlaces covalentes</h4>
 
 Los átomos no metálicos frecuentemente forman enlaces covalentes con otros átomos no metálicos. Por ejemplo, la molécula de hidrógeno, H2, contiene un enlace covalente entre sus dos átomos de hidrógeno. Dos átomos de hidrógeno separados con una energía potencial particular se acercan entre sí, sus orbitales de valencia (1s) comienzan a solaparse. Los electrones individuales en cada átomo de hidrógeno entonces interactúan con ambos núcleos atómicos, ocupando el espacio alrededor de ambos átomos. La fuerte atracción de cada electrón compartido hacia ambos núcleos estabiliza el sistema, y la energía potencial disminuye a medida que disminuye la distancia de enlace. Si los átomos continúan acercándose entre sí, las cargas positivas en los dos núcleos comienzan a repelerse entre sí, y la energía potencial aumenta. La longitud de enlace es determinada por la distancia a la cual se consigue la menor energía potencial.
 
 Es esencial recordar que se debe añadir energía para romper los enlaces químicos (un proceso endotérmico), mientras que la formación de los enlaces químicos libera energía (un proceso exotérmico). En el caso de H2, el enlace covalente es muy fuerte; se debe añadir una gran cantidad de energía, 436 kJ, para romper los enlaces en un mol de moléculas de hidrógeno y hacer que los átomos se separen:
 
 <img alt="Eq1" src="https://www.jovecore.com/CDNSource/lessons/1059/1059_Equations-01.png" style="height: 40px;" />
 
 Por el contrario, la misma cantidad de energía se libera cuando se forma un mol de moléculas de H2 a partir de dos moles de átomos de H:
 
 <img alt="Eq2" src="https://www.jovecore.com/CDNSource/lessons/1059/1059_Equations-02.png" style="height: 40px;" />
 
 <h4>Estructuras de Lewis</h4>
 
 Los símbolos de Lewis se pueden utilizar para indicar la formación de los enlaces covalentes, los cuales se muestran en las estructuras de Lewis, dibujos que describen los enlaces en moléculas e iones poliatómicos. Por ejemplo, cuando dos átomos de cloro forman una molécula de cloro, comparten un par de electrones:
 
 <img alt="Figura1" src="https://www.jovecore.com/CDNSource/lessons/1059/1059Image_1.svg" style="height: 90px;" />
 
 La estructura de Lewis indica que cada átomo de Cl tiene tres pares de electrones que no se utilizan en el enlace (llamados pares solitarios) y un par compartido de electrones (escrito entre los átomos). A veces se utiliza un guión (o línea) para indicar un par compartido de electrones: Cl—Cl.
 
 <ul>
 <li>Un solo par compartido de electrones se denomina enlace único. Cada átomo de Cl interactúa con ocho electrones de valencia: Los seis de los pares solitarios y los dos del enlace único.</li>
 <li>Sin embargo, un par de átomos puede necesitar compartir más de un par de electrones para lograr el octeto requerido. Un enlace doble se forma cuando dos pares de electrones se comparten entre un par de átomos, como entre los átomos de carbono y oxígeno en CH2O (formaldehido) y entre los dos átomos de carbono en C2H4 (etileno).  
 <img alt="Figura2" src="https://www.jovecore.com/CDNSource/lessons/1059/1059Image_2.svg" style="height: 120px;" /></li>
 <li>Un enlace triple se forma cuando tres pares de electrones son compartidos por un par de átomos, como en el monóxido de carbono (CO) y el ion cianuro (CN–). 
 <img alt="Figura3" src="https://www.jovecore.com/CDNSource/lessons/1059/1059Image_3.svg" style="height: 90px;" /></li>
 </ul>
 
 La tabla periódica se puede utilizar para predecir el número de electrones de valencia en un átomo y el número de enlaces que se formarán para alcanzar un octeto. Los elementos del Grupo 18 como el Argón y el Helio han completado su configuración electrónica y, por lo tanto, rara vez participan en los enlaces químicos. Sin embargo, los átomos del grupo 17, como el bromo o el yodo, sólo necesitan un electrón para alcanzar el octeto. Por lo tanto, los átomos que pertenecen al grupo 17 pueden formar un enlace covalente único. Los átomos del grupo 16 necesitan 2 electrones para alcanzar un octeto; por lo tanto, pueden formar dos enlaces covalentes. Del mismo modo, el carbono que pertenece al grupo 14 necesita 4 electrones más para alcanzar un octeto; por lo tanto, el carbono puede formar cuatro enlaces covalentes. 
 
 Este texto es adaptado de<a href="https://openstax.org/books/chemistry-2e/pages/7-2-covalent-bonding">Openstax, Química 2e, Sección 7.2: Enlaces Covalentes</a> y <a href="https://openstax.org/books/chemistry-2e/pages/7-3-lewis-symbols-and-structures">Openstax, Química 2e, Sección 7.3: Símbolos y Estructuras de Lewis.</a>

Enlaces Covalentes y Estructuras de Lewis

En comparación con los enlaces iónicos, que resultan de la transferencia de electrones entre átomos metálicos y no metálicos, los enlaces covalentes son el resultado de la mutua ...

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Chemical Bonding: Basic Concepts

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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

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Shared Electrons in Covalent Molecules | JoVE K12

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Shared Electrons in Covalent Molecules

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