Name: electron-pair reactions in acid-base chemistry
Uploaded: 2020-09-24T15:06:51.000Z
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Lewis Acids and Bases

El modelo de Brønsted-Lowry define los ácidos y las bases en términos de protones, donde los ácidos son donantes de protones y las bases aceptores de protones. Por el contrario, el modelo de Lewis define los ácidos y las bases en términos de pares de electrones, donde los ácidos de Lewis son son aceptores de pares de electrones y las bases de Lewis son donantes de pares de electrones. En un ácido de Brønsted, como el ácido acético, el hidrógeno también puede actuar como un ácido de Lewis porque tiene un orbital vacío para aceptar electrones donados de una base, como el agua, que actúa como una base de Lewis.La ventaja del modelo de Lewis es que permite a los científicos clasificar una mayor cantidad de compuestos como ácidos, incluidos los que no tienen protones ionizables. Por ejemplo, el trifluoruro de boro no se puede clasificar como ácido por el modelo de Brønsted-Lowry porque no contiene hidrógeno. Sin embargo, el trifluoruro de boro posee un octeto incompleto con un orbital vacío que puede aceptar un par de electrones de una base de Lewis, como el amoníaco y, por lo tanto, puede actuar como un ácido de Lewis.El producto resultante formado por tales reacciones ácido-base de Lewis se denomina aducto ácido-basico de Lewis. Algunas moléculas, como el dióxido de carbono, pueden reorganizar sus electrones para actuar como un ácido de Lewis. Por ejemplo, en la reacción entre el agua y el dióxido de carbono, un par de electrones se mueve desde el enlace pi carbono-oxígeno hasta el oxígeno terminal del dióxido de carbono.El orbital vacío resultante en el átomo de carbono le permite aceptar el par de electrones de una molécula de agua y actuar como un ácido de Lewis. A medida que la molécula de agua dona el par de electrones, actúa como una base de Lewis. En una reordenación adicional, se transfiere un protón del oxígeno del agua al oxígeno terminal del dióxido de carbono, lo que da como resultado la formación del aducto de ácido carbónico.Los cationes metálicos pequeños, como el yoduro de aluminio o Al pueden volver a adquirir pares de electrones y actuar como ácidos de Lewis. Por ejemplo, el Al acepta pares solitarios de electrones del agua y forma iones de hexaacualuminio. Aquí, las moléculas de agua donan pares de electrones y actúan como una base de Lewis.

Ácidos y bases de lewis

En 1923, G. N. Lewis propuso una definición generalizada del comportamiento ácido-base en la cual los ácidos y las bases se identifican por su capacidad de aceptar o donar un par de electrones y formar un enlace covalente coordinado.
 Un enlace covalente coordinado (o enlace dativo) ocurre cuando uno de los átomos en el enlace proporciona ambos electrones del enlace. Por ejemplo, un enlace covalente coordinado ocurre cuando una molécula de agua se combina con un ion hidrógeno para formar un ion hidronio. Un enlace covalente coordinado también ocurre cuando una molécula de amoníaco se combina con un ion de hidrógeno para formar un ion de amonio. Estas dos ecuaciones se muestran aquí.
 <img alt="Image1" src="/files/ftp_upload/11407/11407_Image_1.svg" style="height: 175px;" />
 Las reacciones que implican la formación de enlaces covalentes coordinados se clasifican como química ácido-base de Lewis. La especie que dona el par de electrones que componen el enlace es una base de Lewis, la especie que acepta el par de electrones es un ácido de Lewis, y el producto de la reacción es un aducto ácido-base de Lewis. Como ilustran los dos ejemplos anteriores, las reacciones ácido-base de Brønsted-Lowry representan una subcategoría de reacciones ácidas de Lewis, específicamente aquellas en las que la especie ácida es H+. A continuación se describen algunos ejemplos que incluyen otros ácidos y bases de Lewis.
 El átomo de boro en el trifluoruro de boro, BF3, tiene sólo seis electrones en su capa de valencia. Al estar por debajo del octeto preferido, BF3 es un ácido de Lewis muy bueno y reacciona con muchas bases de Lewis; un ion flúor es la base de Lewis en esta reacción, donando uno de sus pares solitarios:
 <img alt="Image2" src="/files/ftp_upload/11407/11407_Image_2.svg" style="height: 150px;" />
 En la siguiente reacción, cada una de las dos moléculas de amoníaco, bases de Lewis, dona un par de electrones a un ión de plata, el ácido de Lewis:
 
 <img alt="Image3" src="/files/ftp_upload/11407/11407_Image_3.svg" style="height: 150px;" />
 Los óxidos no metálicos actúan como ácidos de Lewis y reaccionan con iones óxido, bases de Lewis, para formar oxianiones:
 <img alt="Image4" src="/files/ftp_upload/11407/11407_Image_4.svg" style="height: 150px;" />
 Muchas reacciones ácido-base de Lewis son reacciones de desplazamiento en las cuales una base de Lewis desplaza a otra base de Lewis de un aducto ácido-base, o en las cuales un ácido de Lewis desplaza a otro ácido de Lewis:
 <img alt="Image5" src="/files/ftp_upload/11407/11407_Image_5.svg" style="height: 350px;" />
 Otro tipo de química ácido-base de Lewis implica la formación de un ión complejo (o un complejo de coordinación) que comprende un átomo central, típicamente un catión de metal de transición, rodeado de iones o moléculas llamadas ligandos. Estos ligandos pueden ser moléculas neutras, como el H2O o NH3, o iones, como el CN− o el OH&8722;. A menudo, los ligandos actúan como bases de Lewis, donando un par de electrones al átomo central.
 Este texto está adaptado de<a href="https://openstax.org/books/chemistry-2e/pages/15-2-lewis-acids-and-bases"> Openstax, Química 2e, Sección 15.2: Ácidos y bases de Lewis</a>.

Ácidos y Bases de Lewis

En 1923, G. N. Lewis propuso una definición generalizada de comportamiento ácido-base en la que los ácidos y las bases se identifican por su capacidad de aceptar o ...

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electron-pair reactions in acid-base chemistry

Electron-pair reactions explain Lewis acid-base chemistry. In 1923, G. N. Lewis described acids and bases by how they accept or donate a pair of electrons. When that electron pair forms a coordinate covalent bond, one atom supplies both bonding electrons. This bond is also called a dative bond.

A simple example is the reaction of water with a hydrogen ion to make a hydronium ion. Ammonia can also combine with a hydrogen ion to form an ammonium ion. In both cases, the reaction creates a Lewis acid-base adduct. The electron-pair donor is the Lewis base, and the electron-pair acceptor is the Lewis acid.

Brønsted-Lowry acid-base reactions fit inside Lewis acid chemistry. In those reactions, the acid species is H+. Other Lewis acids show the same electron-pair pattern in different ways. Boron trifluoride, BF3, is a strong Lewis acid because boron has only six electrons in its valence shell and still needs an octet. A fluoride ion can donate a lone pair to BF3 and form an adduct.

Lewis acids can also bind more than one base. In one example, two ammonia molecules each donate a pair of electrons to a silver ion, which acts as the Lewis acid. Nonmetal oxides also behave as Lewis acids and react with oxide ions, which are Lewis bases, to form oxyanions. Many Lewis acid-base reactions are displacement reactions. In these reactions, one Lewis base can replace another base in an adduct, or one Lewis acid can replace another acid.

Lewis acid-base chemistry also helps form complex ions, or coordination complexes. These complexes have a central atom, often a transition metal cation, surrounded by ligands. Ligands can be neutral molecules such as H2O or NH3, or ions such as CN- or OH-. In many cases, the ligands act as Lewis bases and donate electron pairs to the central atom.

Electron-Pair Reactions in Acid-Base Chemistry

electron-pair-reactions-in-acid-base-chemistry

bronsted-lowry acids and bases

Proton Transfer in Acids and Bases

Acids and bases can be understood through proton transfer. Early scientists described acids by their sour taste, their effect on dyes, and their reaction with limestone, which releases carbon dioxide. Robert Boyle also noted that acid solutions can lose these properties after contact with alkali, or base, solutions. The modern view developed step by step. Humphry Davy showed that hydrogen is the essential part of acids. Joseph Louis Gay-Lussac then described acids as substances that neutralize...

water: a bronsted-lowry acid and base

Water as an Amphiprotic Substance

Water can act as both a Brønsted-Lowry acid and a Brønsted-Lowry base. When hydrogen fluoride dissolves in water, it donates protons to water. This is acid ionization. The products are hydronium ions and fluoride ions. Water can also accept protons from another substance. When pyridine dissolves in water, it accepts a proton from water molecules. This is base ionization. The products are hydroxide ions and pyridinium ions. These reactions show that water can play either role. A substance that...

ph scale

pH and pOH in Water

pH and pOH show how hydronium and hydroxide ions behave in water and other aqueous solutions. These ion concentrations are linked by the ion product of water, K w , so when one goes up, the other goes down. Because ion concentrations can differ by huge amounts, a logarithmic scale is used to describe them. pH is defined from the hydronium ion concentration, [H3O+], and pOH is the matching p-function for hydroxide ion concentration. These two values are also connected through the K w...

relative strengths of conjugate acid-base pairs

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How Ka and Kb Show Acid-Base Strength

Brønsted-Lowry acid-base chemistry is based on proton transfer. That idea helps show how the strength of an acid is linked to the strength of its conjugate base, and vice versa. Acid and base strength is measured with ionization constants, Ka and Kb, which tell how far an ionization reaction goes. For the conjugate pair HA and A−, the acid ionizes in water and the base reaction can be written for the partner species. When the two reactions are added, they give the autoionization of water. The...

strong acid and base solutions

Strong Acids, Strong Bases, and pH

Strong acids and strong bases change the pH of a solution by dissociating completely in water. A strong acid produces hydronium ions, which are H3O+ ions, and a strong base produces hydroxide ions, OH-. Because the dissociation is complete, the ion concentration can be found from the starting concentration and the molar ratio. For example, 0.20 M hydrobromic acid dissociates fully in water. It produces 0.20 M hydronium ions and 0.20 M bromide ions. A 0.015 M KOH solution also dissociates...

weak acid solutions

Weak Acid Equilibrium in Water

Weak acid equilibrium in water explains why many acids only partially react when dissolved. Most weak acid molecules stay in their original form, and only a small amount of hydronium ions forms. Weak acids are common in nature. They help create the tangy taste of citrus fruits, the stinging sensation from insect bites, and the smells linked to body odor. Acetic acid is a familiar weak acid. It is the main ingredient in vinegar. Its equation uses a double arrow to show that the reaction with...

weak base solutions

Weak Base pH and Kb Calculations

Weak base pH and Kb calculations show how partially reacting bases change water chemistry. Weak bases produce hydroxide ions when they dissolve by reacting with water molecules. Because the reaction only happens partly, these compounds are classified as weak bases. Weak bases are common in nature and in technology. Ammonia is one major example. Global production of ammonia is typically well over 100 metric tons each year. It is used as an agricultural fertilizer, a starting material for...

mixtures of acids

pH of Acid Mixtures

The pH of acid mixtures depends on the acids present, their dissociation constants, and their starting concentrations. When a solution contains two different acids, the pH can be found by choosing the method that fits their relative strengths. This makes acid mixtures a useful case for applying acid-base rules in chemistry. In a mixture of a strong acid and a weak acid, the strong acid dissociates completely and supplies almost all of the hydronium ions in the solution. The weak acid...

ions as acids and bases

Salt Hydrolysis and pH of Solutions

Salt hydrolysis explains why some salt solutions are acidic, basic, or neutral. Salts are ionic compounds made of cations and anions. In water, either ion may react with water through an acid or base ionization reaction. The pH of the solution depends on the relative acid-base strength of the ions in the salt. Ammonium chloride is a good example of an acidic salt. When it dissolves, it separates into ammonium ions and chloride ions. The ammonium ion is the conjugate acid of ammonia, so it can...

determining the ph of salt solutions

Salt Hydrolysis and pH in Solutions

Salt hydrolysis and pH in solutions depend on the ions released when a salt dissolves. Some ions react with water, while others do not. The balance between the cation and anion helps show whether the solution becomes acidic, basic, or neutral. Salts with pH-neutral anions and cations that produce hydronium ions give acidic solutions with pH below 7. Ammonium nitrate is one example. In NH4NO3, NO3− does not react with water, but NH4+ forms hydronium ions, so the solution is acidic. Other salts...

acid/base strengths and dissociation constants

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Acid and Base Strength in Water

Acid and base strength in water depends on how much a substance ionizes after it dissolves. If ionization is nearly complete, the substance is called strong. If only a small amount ionizes, it is called weak. Weak acids and weak bases are much more common than strong ones. The most common strong acids are HClO4, HCl, HBr, HI, HNO3, and H2SO4. The most common strong bases are LiOH, NaOH, KOH, Ca(OH)2, Sr(OH)2, and Ba(OH)2. These strong acids and bases ionize essentially completely in water. <br...

polyprotic acids

Stepwise Ionization of Polyprotic Acids

Polyprotic acids lose hydrogen ions in steps. The number of ionizable protons in a molecule determines whether an acid is monoprotic, diprotic, or triprotic. This video also shows that polyprotic bases can accept more than one proton. Monoprotic acids have one ionizable hydrogen atom per molecule. Examples include HCl, HNO3, and HCN. Their reactions with water form hydronium ions and the matching anion. Acetic acid, CH3CO2H, is also monoprotic because only the hydrogen in the carboxyl group...

acid strength and molecular structure

How Bonding Affects Acid Strength

Acid strength depends on bonding, electronegativity, and oxidation number. These factors help explain why some hydrogen compounds act as strong acids, while others behave as weak acids or bases. For binary acids, which are hydrogen compounds with one other nonmetal, acidity increases as the H-A bond gets weaker down a group. That is why the order for group 17 is HF &lt; HCl &lt; HBr &lt; HI, and for group 16 it is H2O &lt; H2S &lt; H2Se &lt; H2Te. Across a period, acid strength increases as...

lewis acids and bases

Electron-pair reactions explain Lewis acid-base chemistry. In 1923, G. N. Lewis described acids and bases by how they accept or donate a pair of electrons. When that electron pair forms a coordinate covalent bond, one atom supplies both bonding electrons. This bond is also called a dative bond. A simple example is the reaction of water with a hydrogen ion to make a hydronium ion. Ammonia can also combine with a hydrogen ion to form an ammonium ion. In both cases, the reaction creates a Lewis...

key terms and concepts

Learn how Lewis acids and bases use electron pairs to form coordinate covalent bonds, adducts, complex ions, and examples like BF3, Ag+, and hydronium.

Electron-Pair Reactions in Acid-Base Chemistry | JoVE K12

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Electron-Pair Reactions in Acid-Base Chemistry

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