Name: electron-pair reactions in acid-base chemistry
Uploaded: 2020-09-24T15:06:51.000Z
Duration: 2 min 33 s
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Lewis Acids and Bases

Das Brønsted-Lowry-Modell definiert Säuren und Basen in Form von Protonen, wobei Säuren Protonendonatoren und Basen Protonenakzeptoren sind. Im Gegensatz dazu definiert das Lewis-Modell Säuren und Basen in Form von Elektronenpaaren, wobei Lewis-Säuren Elektronenpaarakzeptoren und Lewis-Basen Elektronenpaardonatoren sind. 
In einer Brønsted-Säure kann der Wasserstoff wie Essigsäure auch als Lewis-Säure wirken, da er ein leeres Orbital hat, um Elektronen aufzunehmen, die von einer Base wie Wasser abgegeben werden, die als Lewis-Base fungiert.
Der Vorteil des Lewis-Modells besteht darin, dass es den Wissenschaftlern ermöglicht, eine größere Anzahl von Verbindungen als Säuren zu klassifizieren – einschließlich solcher, die keine ionisierbaren Protonen enthalten. Zum Beispiel kann Bortrifluorid nach dem Brønsted-Lowry-Modell nicht als Säure eingestuft werden, da es keinen Wasserstoff enthält.
Bortrifluorid besitzt jedoch ein unvollständiges Oktett mit einem leeren Orbital, das ein Elektronenpaar von einer Lewis-Base, wie z. B. Ammoniak, aufnehmen und daher als Lewis-Säure wirken kann.
Das resultierende Produkt, das durch solche Lewis-Säure-Base-Reaktionen gebildet wird, wird als Lewis-Säure-Base-Addukt bezeichnet.
Einige Moleküle, wie Kohlendioxid, können ihre Elektronen so umanordnen, dass sie als Lewis-Säure wirken.
Bei der Reaktion zwischen Wasser und Kohlendioxid wandert beispielsweise ein Elektronenpaar von der Kohlenstoff-Sauerstoff-Pi-Bindung zum terminalen Sauerstoff des Kohlendioxids.
Das resultierende leere Orbital am Kohlenstoffatom ermöglicht es ihm, das Elektronenpaar von einem Wassermolekül aufzunehmen und als Lewis-Säure zu wirken. Wenn das Wassermolekül das Elektronenpaar abgibt, fungiert es als Lewis-Base.
Bei weiterer Umlagerung wird ein Proton vom Wassersauerstoff auf den terminalen Sauerstoff des Kohlendioxids übertragen, wodurch das Kohlensäureaddukt gebildet wird.
Kleine Metallkationen wie Al(III) können Elektronenpaare wieder aufnehmen und als Lewis-Säuren wirken. Zum Beispiel nimmt Al(III) einsame Elektronenpaare aus Wasser auf und bildet Hexaaquaaluminium-Ionen. Hier geben Wassermoleküle Elektronenpaare ab und fungieren als Lewis-Base.

lewis-säuren und -basen

Im Jahr 1923 schlug G. N. Lewis eine verallgemeinerte Definition des Säure-Base-Verhaltens vor, in der Säuren und Basen durch ihre Fähigkeit identifiziert werden, ein Elektronenpaar aufzunehmen oder abzugeben und eine koordinative kovalente Bindung zu bilden. Eine koordinative kovalente Bindung (oder dative Bindung) entsteht, wenn eines der Atome in der Bindung beide Bindungselektronen bereitstellt. Eine koordinative kovalente Bindung entsteht beispielsweise, wenn sich ein Wassermolekül mit einem Wasserstoffion zu einem Hydroniumion verbindet. Eine koordinative kovalente Bindung entsteht auch, wenn sich ein Ammoniakmolekül mit einem Wasserstoffion zu einem Ammoniumion verbindet. Beide Gleichungen werden hier gezeigt. <img alt="Image1" src="/files/ftp_upload/11407/11407_Image_1.svg" style="height: 175px;"> Reaktionen, die die Bildung koordinativer kovalenter Bindungen beinhalten, werden als Lewis-Säure-Base-Chemie klassifiziert. Die Spezies, die das Elektronenpaar spendet, aus dem die Bindung besteht, ist eine Lewis-Base, die Spezies, die das Elektronenpaar aufnimmt, ist eine Lewis-Säure, und das Produkt der Reaktion ist ein Lewis-Säure-Base-Addukt. Wie die beiden obigen Beispiele veranschaulichen, stellen Brønsted-Lowry-Säure-Base-Reaktionen eine Unterkategorie der Lewis-Säure-Reaktionen dar, insbesondere solche, bei denen die Säurespezies H+ ist. Nachfolgend werden einige Beispiele für andere Lewis-Säuren und -Basen beschrieben. Das Boratom in Bortrifluorid, BF3, hat nur sechs Elektronen in seiner Valenzschale. Da BF3 nicht über das bevorzugte Oktett verfügt, ist es eine sehr gute Lewis-Säure und reagiert mit vielen Lewis-Basen; Ein Fluoridion ist in dieser Reaktion die Lewis-Base und spendet eines seiner freien Elektronenpaare: <img alt="Image2" src="/files/ftp_upload/11407/11407_Image_2.svg" style="height: 150px;"> Bei der folgenden Reaktion gibt jedes der beiden Ammoniakmoleküle, Lewis-Basen, ein Elektronenpaar an ein Silberion, die Lewis-Säure, ab: <img alt="Image3" src="/files/ftp_upload/11407/11407_Image_3.svg" style="height: 150px;"> Nichtmetalloxide wirken als Lewis-Säuren und reagieren mit Oxidionen, Lewis-Basen, unter Bildung von Oxyanionen: <img alt="Image4" src="/files/ftp_upload/11407/11407_Image_4.svg" style="height: 150px;"> Bei vielen Lewis-Säure-Base-Reaktionen handelt es sich um Verdrängungsreaktionen, bei denen eine Lewis-Base eine andere Lewis-Base aus einem Säure-Base-Addukt verdrängt oder bei denen eine Lewis-Säure eine andere Lewis-Säure verdrängt: <img alt="Image5" src="/files/ftp_upload/11407/11407_Image_5.svg" style="height: 350px;"> Eine andere Art der Lewis-Säure-Base-Chemie beinhaltet die Bildung eines komplexen Ions (oder eines Koordinationskomplexes), das ein Zentralatom, typischerweise ein Übergangsmetallkation, umfasst, das von Ionen oder Molekülen, sogenannten Liganden, umgeben ist. Diese Liganden können neutrale Moleküle wie H2O oder NH3 oder Ionen wie CN– oder OH– sein. Oft fungieren die Liganden als Lewis-Basen und geben ein Elektronenpaar an das Zentralatom ab. Dieser Text wurde angepasst von<a href="https://openstax.org/books/chemistry-2e/pages/15-2-lewis-acids-and-bases"> Openstax, Chemistry 2e, Section 15.2: Lewis Acids and Bases</a>.

Lewis-Säuren und -Basen

Im Jahr 1923 schlug G. N. Lewis eine verallgemeinerte Definition des Säure-Base-Verhaltens vor, in der Säuren und Basen durch ihre Fähigkeit identifiziert werden, ...

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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...

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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

relative-strengths-of-conjugate-acid-base-pairs

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...

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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...

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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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